Gen Chem Exam 3 Flashcards
(220 cards)
1
Q
What undergoes combustion during the welding of metals?
A
A mixture of acetylene and oxygen
2
Q
What is more practical and convenient to measure compared to moles?
A
Mass
3
Q
Limiting Reactant
A
- The substance that is used up first
- The substance that limits the amount of product that can form.
4
Q
Excess Reactant
A
The reactants that do not completely react and are left over at the end of the reaction.
5
Q
What reactant yields the smallest amount of product?
A
Limiting reactant
6
Q
Theoretical Yield
A
The maximum amount of product, which is calculated using the balanced equation and the limiting reactant.
7
Q
Actual Yield
A
The amount of product experimentally obtained or collected during the experiment.
8
Q
Percent Yield
A
The ratio of actual yield to theoretical yield.
9
Q
Percent Yield Formula
A
Actual Yield/Theoretical Yield x 100%
10
Q
During what 3 scenarios does energy change occur?
A
- When reactants interact
- When bonds break apart
- When products form
11
Q
Heat of Reaction / Enthalpy Change
A
The difference between the enthalpy of the products and enthalpy of the reactants
12
Q
Exothermic Reaction
A
- Heat is released
- The energy of the products is less than the energy of the reactants
- Heat is a product
13
Q
Endothermic Reactions
A
- Heat is absorbed
- The energy of the products is greater than the energy of the reactants
- Heat is a reactant
14
Q
Cold packs being activated are an example of a ___ reaction.
A
endothermic
15
Q
How many elements in the periodic table exist as gases at room temperature conditions?
A
Only a few
16
Q
What elements are considered gases at room temperature conditions?
A
- All noble gases
- Some diatomic molecules (H2, N2, O2, F2, and Cl2)
- Oxides of the nonmetals
(CO, CO2, NO, NO2, SO2, and SO3)
17
Q
What are the 7 properties of gases?
A
- Indefinite shape and volume
- Takes the same shape and volume as its container
- Particles are far apart
- Particles are moving quickly in random directions
- Low densities
- Compressible
- Form homogenous mixtures
18
Q
Kinetic Molecular Theory of Gases
A
- Gases consist of tiny particles moving randomly at high velocities
- Most of the volume of a gas is empty space.
- There are no attractive or repulsive forces between gas particles
- Gas particles are constantly moving in straight paths.
- The average kinetic energy of the gas particles is proportional to the kelvin temperature of the sample.
19
Q
What causes gas pressure to arise?
A
Collisions exerted by gas particles against the container walls.
20
Q
What four units are used to measure pressure?
A
- Pascal (Pa)
- mmHg
- Torr
- Atmosphere (atm)
21
Q
Atmosphere
A
A mixture of gases that cover the earth
22
Q
Atmospheric Pressure
A
The pressure exerted by a column of this air mixture towards the surface of the Earth.
23
Q
When is atmospheric pressure (atm) at 1?
A
Sea level
24
Q
How many mmHg and Torr are in 1 atm?
A
760
25
Barometer
Measures the pressure exerted by the gases in the atmosphere.
26
What causes atmospheric pressure to change?
1. Weather
2. Altitude
27
What weather causes the atmospheric pressure to increase?
Hot and sunny
28
What weather causes the atmospheric pressure to decrease?
Rainy
29
Atmospheric pressure decreases as ___ increases.
Altitude
30
What are the 4 variables that define the physical properties of a gas?
1. Pressure (P)
2. Temperature (T)
3. Volume (V)
4. Amount of gas moles (n)
31
What are the 4 gas laws?
1. Boyle's Law
2. Charles's Law
3. Gay-Lussac's Law
4. Avogadro's Law
32
Ideal Gas Law Definition
A gas whose behavior follows the 4 gas laws.
33
Ideal Gas Law Formula
PV = nRT
34
What does the variable R represent in the ideal gas law formula?
The gas constant (value is given on sheet)
35
What type of volumes do cylinders have?
A fixed volume (can't be greater than 1 atm)
36
What happens to the volume when you blow up a balloon?
The volume will expand until the pressure equalizes with the atmospheric pressure.
37
Combined Gas Law Definition
When more than two properties of gas change at once, two laws can be combined to explain its behavior.
38
Combined Gas Law Equation
(P1 x V1)/T1 x (P2 x V2)/T2
39
Boyle's Law
V decreases as P increases
- n and T are constant
40
Charles's Law
V changes the same to T
- n and P are constant
41
Gay- Lussac's Law
P changes the same to T
- n and V are constant
42
What causes vapor to form?
When liquid molecules with enough kinetic energy break away from the surface of a liquid.
43
What happens when liquid is left in an open container?
It will eventually evaporate
44
What happens when vapor is left in a closed container?
The vapor accumulates and creates pressure called vapor pressure.
45
What is vapor pressure of a liquid dependent on?
Temperature
46
Boiling Point
The temperature at which the vapor pressure (inside container) becomes equal to the external pressure.
47
Water boils at ___ temperatures in the mountains.
lower
48
Avogadro's Law
V changes the same to n
- P and T are constant
49
What are the standard conditions for the gas formulas?
T = Kelvin
V = L
n = Moles
P = atm
50
Why do the gas formulas need to be converted with the same units?
So comparisons can be made among different experiments.
51
What is the standard volume at 1 mole of gas?
22.4 L (molar volume)
52
What does the molar volume mean?
No matter the identity of the gas, the volume is ALWAYS 22.4 L at 1 mole.
53
Dalton's Law of Partial Pressures
The total pressure exerted by a mixture of gases in a container at constant V and T
=
The sum of the pressures of each individual gas in the container
54
Partial Pressure
A gases own pressure in comparison to the total pressure of the gas.
55
Partial Pressure Formula
P total = P1 + P2 + P3 + Pn...
56
What are the different elements within the air we breath?
1. N2 (78%)
2. O2 (21%)
3. other gases (1%)
57
What is the partial pressure exerted by a given gas proportional to?
The mole percentage of that gas in the mixture.
58
What are the 4 IMFs?
1. Ion Dipole
2. Hydrogen Bonding
3. Dipole-Dipole
4. Dispersion
59
Aqueous Solution (aq)
Dissolved in water
60
Solutions
Homogenous (completely mixed) mixtures of 2+ substances
61
What are the two components of solutions?
1. Solvents
2. Solutes
62
Solvent
Present in a larger amount
- major component of the slution
63
Solute
Present in a smaller amount
- spread evenly throughout the solution
64
___ dissolves into ___ to form a solution.
Solute, solvent
65
What physical states can a solute have?
Liquid, gas or solid
66
How can solutes be seperated?
Evaporation
- not by filtration
67
Are solutes visible?
No but they can add a color to the solution
68
What causes the solute to dissolve into the solvent?
Molecules within the solvent are capable of forming sufficient attractions with the solute particles to make them dissolve
69
Water (H2O)
- A polar molecule due to O-H bonds
70
What is one of the most common solvents in nature?
Water
71
What do solute molecules form with water?
Intermolecular forces
72
What takes place in water media?
The vast majority of chemical reactions (biological reactions)
73
Aqueous Reactions
Reactions that take place in water media
74
What combinations causes solutions to form?
Polar - Polar
Nonpolar - Nonpolar
75
What are the 3 common nonpolar solvents?
1. Hexane (C6H12)
2. Toluene (C7H8)
3. Carbon Tetrachloride (CCL4)
76
Do nonpolar solutes dissolve in polar solvents?
NO
- Nonpolar outweighs the polar
77
What are the 2 examples of nonpolar solutes?
1. Hydrocarbons, oils, and fats
2. Compounds that can only form dispersion forces (cancel out)
78
Hydrocarbons
Compounds with only hydrogen and carbon
79
What happens when there is an absence of attractions?
A solution cannot form due to insufficient energy levels.
80
Nonpolar molecular compounds do not dissolve in....
Water
81
What are the 3 most common polar solvents?
1. Water (H2O)
2. Acetone ((CH3)2C=O)
3. Ethanol (CH3CH2OH)
82
What are examples of polar solutes?
1. Ionic compounds or salts
2. Compounds that can form dipole-dipole interactions and/or Hydrogen Bonding
83
Within a lewis structure, hydrogen bonds are shown as...
dashed lines
84
Are polar molecular compounds, such as menthol, soluble in water?
Yes
- There is a polar OH- group present to form H bonds in water
85
What are 2 examples of polar molecular compounds?
1. Small polar molecules
2. Water soluble vitamins
86
Ionic Solutes
Form ion dipole interactions with water molecules.
87
What is an example of an ionic solute?
NaCl crystals
88
What happens during the ionic solute of NaCl crystals?
1. Na+ and Cl- ions are pulled out into the solution and get dissolved
89
What does hydrated mean?
Surrounded by water
(water molecules surround themselves around the positive and negative ions)
90
Dissociate
A solid dissolves in water
- becoming hydrated as an ion
91
How are solutes in a solution classified?
1. Electrolytes
2. Non-electrolytes
92
Dissolved ions play an important role in...
Maintaining the proper function of the nervous system, cells, and organs.
93
How can dissolved ions such as K, Cl, and bicarbonate be measured?
In a blood test
94
Electrolytes
Compounds that dissociate to produce ions when dissolved in water.
95
What do electrolytes conduct?
Electricity due to the presence of ions
96
What are examples of electrolytes?
1. Water soluble ionic compounds
Ex) NaCl, KBr, Ca(NO3)2, LiOH, CuSO4
2. Molecular compounds that are acids
Ex) HCl, H2SO4, COOH
97
How are electrolytes classified?
1. Strong
2. Weak
98
Strong Electrolytes
Completely dissolve in water to produce ions.
- Distinguished by a forward pointing arrow
99
What are examples of strong electrolytes?
1. Water soluble ionic compounds
Ex) NaCl, KBr, Ca(NO3)2, LiOH, CuSO4
2. Molecular compounds that are STRONG acids
Ex) HCL, H2SO4
100
Weak Electrolytes
Dissolve only slightly in water to produce few ions.
- Mixture of ions and intact molecules
101
A formula with weak electrolytes is displayed by...
A double arrow
= does not completely dissolve
102
Nonelectrolytes
Compounds that do not dissolve in water to produce ions.
- only intact molecules are found
103
Non electrolytes do not...
conduct electricity due to the absence of ions
104
What are examples of nonelectrolytes?
Molecular compounds that are not acids like carbon containing compounds.
- Ethanol, sucrose, urea
105
Solubility
The maximum amount of solute that dissolved in a specific amount of solvent at a specific heat.
106
How is solubility expressed?
As grams of solute in 100 grams of solvent (water)
- (g solute)/(100 g water)
107
What is the solubility of a given solute dependent on?
Temperature
108
Unsaturated Solutions
Contain less than the maximum amount of solute (solubility)
109
What are two characteristics of unsaturated solutions?
1. No undissolved solute found at the bottom of the container.
2. If more solute is added, it will dissolve.
- No noticeable materials within the solution.
110
Saturated Solutions
Contains the maximum amount of solute that can dissolve at a specific temperature
111
What are 3 characteristics of saturated solutions?
1. Have undissolved solute at the bottom of the container
2. If more solute is added, it will NOT dissolve
3. Contain solute that dissolves as well as solute that recrystallizes in an equilibrium process.
- displayed with double arrow (not fully dissolved)
112
What is the relationship between temperature and solubility in solids?
As the temperature increases, the solubility increase
113
What physical state usually becomes more soluble in water when the solution is heated up?
Solids
114
What is the relationship between temperature and solubility in gases?
As the temperature increases, the solubility level decreases
115
Henry's Law
The solubility of a gas in a liquid is directly related to the pressure of the gas above the liquid.
116
What is the relationship between pressure and gas molecules?
As the pressure level increases, more gas molecules dissolve in the liquid.
117
Insoluble Ionic Compound
Will form when the ionic bonds are too strong for the polar molecules to break
=
visually separation within the solution (precipitate)
118
Solubility Guidelines
A set of rules that help determine whether a compound is soluble or not.
(given on exam sheet)
119
Precipitation Reactions
A chemical reaction in which soluble reactants yield an insoluble product that precipitates out of a solution
120
Precipitate
A solid
121
What happens when two soluble solutions are mixed together?
An insoluble solution is created.
122
What type of reactions do precipitation reactions undergo?
Double Reaction
- the cation of one compound partners with the anion of the other compound.
123
Complete Ionic Equation
All soluble reactants (aq) and products are written as dissociated ions.
124
How are solids written in an ionic equation?
Intact since they are no longer dissolved in the solution
125
Net Ionic Equation
Remove the spectator ions from the complete ionic equation.
126
Spectator Ions
Ions that remain unchanged from the reactant to product side
(stays the same)
127
Concentration
How much of a solute is dissolved in a solution
128
Concentration of a solution =
(amount of solute)/
(amount of solution)
129
What units can solute be expressed in?
-grams
-millimeters
-moles
130
What units can a solution be measured in?
- grams
- millimeters
- liters
131
Mass/Volume Percent of a Solute
mass/volume percent
=
(mass of solute)/(volume of solution) X 100
132
Mass Percent of a Solute
mass percent
=
(mass of a solute)/(mass of a solution) X 100
133
Volume Percent of a Solute
Volume percent
=
(volume of solute)/(volume of solution) X 100
134
Molarity of a Solute
Molarity (M)
=
(moles of solute)/(volume of solution)
135
What happens to concentrated solutions before use?
It is diluted
136
What happens to the volume/concentration ratio when a chemical is diluted?
- Volume increases
- Concentration decreases
137
What happens to the Molarity (M) when a solution is diluted?
Molarity (M) decreases because the volume increases
138
What MUST stay the same throughout the dilution process?
Number of moles of solute
139
How are chemicals stored or prepared?
As concentrated solutions
140
What has to be done to concentrated solutions before they are used?
They must be diluted
141
What are common examples of concentrated solutions?
Acids
142
What happens to the amount of concentrated solution when it is diluted?
It is now spread out in a larger volume of water.
143
What is the formula of dilution?
M1 X V1 = M2 X V2
144
What does 1 and 2 represent in the dilution formula?
Initial (1) and Final (2)
145
What are the 3 types of mixtures?
1. Solution
2. Colloid
3. Suspension
146
What do the properties of solutions measure?
Size vs quanitity
147
What type of particles are solutions?
Small particles
(atoms, ions, or small molecules)
148
How do particles settle in a solution?
The particles DO NOT settle
149
How can particles be separated in a solution?
They cannot be separated.
150
What types of particles make up a colloid?
Larger molecules or groups of molecules/ions
151
How do particles settle in a colloid?
The particles DO NOT settle
152
How can particles be separated in a colloid?
By semipermeable membranes.
153
What types of particles are in a suspension?
VERY large particles
- may be visible
154
How do particles in a suspension settle?
Settle rapidly
155
How can particles in a suspension be separated?
By filters
156
What are the two ways to separate mixtures?
- Filters
- Semipermeable Membranes
157
What are examples of suspensions?
- Blood platelets
- Muddy water
- Calamine lotion
- Flour in water
(separates on standing)
158
What type of mixture is heterogenous?
Suspensions
159
What type of mixture is transparent?
Solutions
160
What is osmotic pressure?
A colligative property
161
Colligative Properties
Depends only on the NUMBER of solute particles in the solution.
- Not focused on concentration/identity (amount)
162
What do semipermeable membranes allow to pass through?
Solvent molecules
163
What do semipermeable membranes NOT allow to pass through?
Solute molecules/ions
164
Osmosis
Solvent molecules (water) flow from the lower solute concentration to the higher solute concentration via a semipermeable membrane.
165
What happens to the flow of water in osmosis?
It becomes equal in both directions
166
What happens to the water level of the solutions in osmosis?
The one with the higher level concentration rises.
167
Osmotic Pressure
The pressure that would prevent the flow of additional water into the more concentrated solution.
168
What causes osmotic pressure to increase?
An increase in amount of dissolved molecules in a solution.
169
Reverse Osmosis
An external pressure that is greater than osmotic pressure is applied to a solution, forcing it through the opposite way of the membrane.
170
What happens to the flow of water in reverse osmosis?
It is reversed
- High to low solute concentration
171
What does reverse osmosis leave behind?
The molecules and ions in the solution.
172
What is one example of reverse osmosis?
It is used in removing the salt water from plants in order to obtain pure water.
173
What is the problem with osmosis?
It requires a large amount of energy
174
What causes osmosis to continuously occur in biological systems?
The cell membranes are semi permeable
175
176
What are examples of solutes in the body?
- Blood
- Tissue fluids
- Lymph
- Plasma
177
What do all solutes in the body do?
Exert osmotic pressure
178
What is an example of isotonic solutions in hospitals?
IV solutions
179
What are the 3 types of solutions?
1. Isotonic
2. Hypotonic
3. Hypertonic
180
Isotonic Solutions
Exert the same osmotic pressure as the solute concentration.
181
Hypotonic Solutions
Exert lower pressure than the solute concentration.
182
Hypertonic Solution
Exert higher osmotic pressure than the solute concentration.
183
What can solute concentrations be compared to when examining solution types?
Red blood cells in the body
184
Hemolysis
Water flows into red blood cells, causing the to swell/burst
185
What type of solution does hemolysis include?
Hypotonic
186
Crenation
Water flows out of the red blood cells, causing them to shrink.
187
What type of solution does crenation involve?
Hypertonic
188
Do products or reactants have more energy in Exothermic Reactions?
Reactants (negative energy level)\
189
Do products or reactants have more energy in Endothermic Reactions?
Products (positive energy level)
190
Energy for Chemical Reactions Formula
(Energy of Products) - (Energy of Reactants)
191
Rate of Reaction Formula
(change in concentration of reactant/product)/
(change in time)
192
Collision Theory
Chemical reactions occur through collisions between molecules/atoms
193
What 3 conditions must be met for a chemical reaction to occur?
1. The reactants must collide
2. The reactants must properly align to break and form bonds
3. The collision must provide sufficient energy of activation
194
Activation Energy (Ea)
The minimum amount of energy required to break the bonds between atoms of the reactants
195
What must be overcome for the reactants to turn to products?
Activation energy barrier
196
What 3 factors influence the rate of a chemical reaction?
1. Changes in temperature
2. Changes in reactant amount (concentration)
3. Presence of a catalyst
197
How does increasing the temperature increase reaction rate?
More collisions occur
198
How does increasing the reactant concentration increase reaction rate?
More collisions occur
199
How does adding a catalyst increase reaction rate?
Lowers the energy of activation
200
What increases as temperature increases?
Kinetic energy of the reactant molecules
201
What happens to reaction rates as the reaction proceeds?
It decreases due to reactants getting converted to products
202
How does adding a catalyst speed up rate of reaction?
Provides an alternate pathway that has a lower activation energy.
- More collisions provide sufficient energy for the reactants to form products
203
What does not happen to the catalyst during a chemical reaction
It is not changed or consumed during the reaction
204
Do reactions with a lower or higher activation energy level have faster reaction rates?
Lower
205
Forward Reaction
Reactants to products
- produced until limiting
206
Reverse Reaction
Products to reactants
207
What two reactions occur at the same time?
Forward and Reverse reactions
208
Reversible Reactions
Reactions that occur in both the forward and reverse direction at the same time
209
Are all reactions reversible?
MOST but not all
210
What is a sign that a reaction is reversible?
Two arrows pointing in opposite directions
211
What is the relationship between the rate of forward reaction and rate of reverse reaction?
Rate of Forward Reaction decreases
as
Rate of Reverse Reaction increases
(as time goes on)
212
Equilibrium
Rate of forward reaction
=
Rate of reverse reaction
213
What is the relationship between the concentration of reactants and the concentration of products?
The concentration of reactants decreases
as
the concentration of products increases
214
What happens to the concentrations of reactants at equilibrium?
No longer change with time
- Reactants and products are being used up at the same rate at which they are being formed
215
What is produced at equilibrium?
The same mixture is formed from the reactants and products.
216
Equilibrium Constant
(Concentration of Products)/
(Concentration of Reactants)
217
What do the coefficients represent in the equilibrium constant formula?
Exponents
218
Does the equilibrium constant have units?
No
219
What unit are concentrations given in?
Molarity units (mol/L)
220
