Honors -Final Exam Flashcards
(352 cards)
1
Q
How to calculate neutrons:
A
of Neutrons = Atomic Mass Number - Atomic Number
1
Q
Electron Configuration
A
A shorthand method of writing the location of electrons by sublevels.
1
Q
How do the bulk properties of polymers effect their behavior?
A
Linear structure - thermoplastics (can be melted and thus recycled.)
Branched chains - strengthen the structure of thermoplastics
Loosely Cross Linked Chains - loosely cross linked structures form an elastomers (rubber)
Tightly Cross Linked Chains - thermoset plastics (can not be melted and thus can not be recycled)
A polymer’s architecture affects many of its physical properties including, but not limited to, solution viscosity, melt viscosity, solubility in various solvents, glass transition temperature and the size of individual polymer coils in solution.
1
Q
Combustion
A
Any compound formed from only Carbon an Hydrogen (fuel) added to Oxygen to produce Carbon Dioxide and Water.
1
Q
What happens to an exothermic reaction when temperature is increased?
A
Heat is a product, so the reaction shifts to the left to make more reactant.
1
Q
How to determine if a solution is unsaturated, saturated, or supersaturated.
A
Add more solute and observe the results.
1
Q
Acids
A
Electron acceptors.
Yeild H+ Ions in solution.
H+ donor.
2
Q
Write the formula for Mg²⁺ and PO₄³⁻
A
Using the criss-cross method and subscripts to insure sum of charges is zero: Mg₃(PO₄)₂
2
Q
Ionic Compounds
A
Conduct Electricity when melted or dissolved in water because….
2
Q
What happens to an endothermic reaction when temperature is increased?
A
Heat is a reactant, so the reaction will shift to the right to make more product.
3
Q
The Mole
A
A unit of measure.
1 mole = the atomic mass of an element
and
1 mole = 6.02 E23 atoms, particles, molecules, units, etc.
3
Q
Solid
A
Solid is one of the four fundamental states of matter. It is characterized by structural rigidity and resistance to changes of shape or volume.
4
Q
Orbital
A
Sublevels can be broken down into regions called “orbitals”. An orbital is defined as the most probable location for finding an electron. Each orbital holds 2 electrons.
5
Q
Thermochemistry
A
The study of energy changes during a chemical reaction or change of state (solid, liquid, gas).
6
Q
Ionic Compounds
A
two words, first names cation second names anion. Indicate charge of transition metal cation by Roman Numeral. MgCl₂ = Magnesium Chloride Cr(NO₃)₃ = Chromium(III) Nitrate SnCl₂ = Tin(II) Chloride
7
Q
Calculate the % Yield of solid silver chromate produced in the following reaction:
K₂CrO₄ + 2AgNO₃ → Ag₂CrO₄ + 2KNO₃
In the reaction there was .500g of the limiting reactant AgNO₃. In the actual experiment, .455g of Ag₂CrO₄ was produced.
A
Calculating the Theoretical Yield of Ag₂CrO₄ that was produced:
0.500g AgNO₃ x (1 mole AgNO₃ / 169.9g AgNO₃) x
(1mole Ag₂CrO₄ / 2 mole AgNO₃) x (331.7 g Ag₂CrO₄ / 1 mol Ag₂CrO₄) =0 .488g Ag₂CrO₄.
Find the ratio of the actual yield (.455g) to the theoretical yield (.488g) …
% Yield = (.455g Ag₂CrO₄ ÷ .488g Ag₂CrO₄ ) x 100 =
93.2 %
7
Q
Amphoteric
A
Acid or Base
8
Q
What is the electron configuration of Vanadium?
A
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³
[Ar] 4s² 3d³
9
Q
Intramolecular Forces
A
Attractive forces WITHIN a molecule (Ionic & Covalent Bonds)
These bonds are stronger than intermolecular forces.
9
Q
Single Replacement
A
An element and a compound are located on each side of the equation. The Activity Series must be checked to determine if the reaction will occur. If the reaction occurs, the element on the reactant side switches with the similarly charged ion within the compound on the reactant side.
A + BC –> AC + B
9
Q
Melting
A
Make or become liquefied by heat. (Solid to Liquid)
10
Q
Atomic Number
A
Identifies an element. The number of protons. For a neutral atom, the number of electrons will equal the number of protons. For an ion (charged atom) the number of electrons will not be the same as the atomic number.
10
Q
Reactants
A
Substances being mixed in a chemical reaction. The left side of the equation.
10
Q
Predict the products of and write a balanced equation from the following statement:
H₂O₂ →
A
H₂O₂ → H₂ + O₂
DECOMPOSITION
11
Is the solution unsaturated, saturated, or supersaturated, if I have 60g of Potassium Nitrate in 100g of Water at 30 C?
Supersaturated
13
What is the orbital diagram for Magnesium?
13
Naming Ionic Bonds
State the name of the metal followed by the nonmetal with the ending changed to "ide". If the metal is a transition metal, add a (Roman numeral). AlCl₃ = Aluminum Chloride CuCl₃ = Copper (III) Chloride
13
The three types of radiation and their penetration abilities:
14
Non Polar Covalent
A type of covalent bond between two atoms in which electrons are shared equally.
15
Percent Yeild
A measure of the efficiency of a reaction carried out in the laboratory.
16
Which of the following elements has the largest atomic radius?
A. nitrogen B. oxygen C. fluorine D. neon
N
(least number of protons attracting the electrons on the energy levels)
17
## Footnote
CaCl₂ + K₂CO₃ →
## Footnote
CaCl₂ + K₂CO₃ →CaCO₃ + 2KCl Double Replacement
18
Writing Formulas for Ionic Bonds
Absolute Value Criss Cross Oxidation #'s
18
What is the IUPAC name of
18
## Footnote
Products of an acid/base reaction
Salt + Water
19
Unsaturated
Contains less solute than it can at that temperature.
20
Atomic Mass Number
Also known as the Mass Number or Atomic Mass.
20
Ionic Bond
Metals GIVE electrons, Nonmetals TAKE electrons
20
Coefficients
The large numbers in front of chemical formulas. Represents the number of molecules of the substance.
2Al
21
What element is represented by this orbital diagram?
Carbon
23
Quantum of Energy
Absorbed when an electron moves from its ground state to its excited state.
23
Lone Pairs
Electron pairs which are not involved in bonding. They do however, affect the shape because electron pairs repel other electron pairs.
23
Oxidation Number
Ion Charge - number of electrons transferred to or away from an atom when it becomes an ion.
23
Law of Conservation of Energy
Energy can not be created nor destroyed during a chemical reaction
24
Calculate the volume required to prepare the diluted solution: Given 6.0M NaOH; need 5.0L of 0.10M NaOH.
M1V1 = M2V2
V1 = (M2V2)/(M1)
V1 = 0.083L
25
Predict the products and write a balanced equation from the following word statement:
Methane (CH₄) plus Oxygen Gas forms:
26
The Dissolving Process
1. The solvent surrounds the outer surface of the solute.
2. Individual Solute Ions are "stripped" away from the surface due to their attraction to the solvent particles.
3. This exposes more particles to the solvent and thus the process continues.
27
Intermolecular Forces
Attractive forces BETWEEN molecules.
Van Der Waals or London Dispersion Forces are the weakest type of intermolecular force and hydrogen bonds are the strongest.
27
## Footnote
C₄H₆ + O₂ →
## Footnote
2C₄H₆ + 7O₂ →4CO₂ + 6H₂O Combustion
28
How to determine if a solution is unsaturated, saturated, or supersaturated.
Add more solute and observe the results.
29
Write a chemical equation from the following word equation:
Lithium chloride reacts with magnesium nitrate to form lithium nitrate and magnesium chloride.
## Footnote
2LiCl + Mg(NO₃)₂ --\> 2Li(NO₃) + MgCl₂
30
Describe crosslinking.
A process that makes a polymer more rigid and resistant to heat. The cross links in slime are hydrogen bonds.
30
Write a chemical equation from the following word equation:
Magnesium reacts with oxygen to form magnesium oxide.
## Footnote
2Mg + O₂ --\> 2MgO
30
How can you make things dissolve faster?
1. Agitation (stirring)
2. Temperature
3. Increase Surface Area (Grind or Crush)
30
Is the solution unsaturated, saturated, or supersaturated, if I have 60g of Potassium Nitrate in 100g of Water at 30 C?
Supersaturated
30
Specific Heat
the amount of energy in the form of heat, calories or joules, required to raise the temperature of 1 gram of a substance 1°C
31
What is the solubility of Potassium Nitrate at 50 C?
86 g Potassium Nitrate
32
In the equation:
4Al + 3O₂→ 2Al₂O₃,
how many moles of Al₂O₃will be produced if there are 3.75 moles of O₂?
3. 75 moles O₂ x (2 moles Al₂O₃/ 3 moles O₂) =
2. 50 moles Al₂O₃
33
## Footnote
Subscripts
## Footnote
Small numbers to the lower right of chemical symbols. Represent the number of atoms of each element in a molecule.
0₂
34
Photon
Energy given off in the form of light by an excited electron. Otherwise known as a "quanta of light".
35
Ions
Atoms or groups of atoms with a charge. To have a charge an atom must have gained or lost electrons. If an atom gains electrons it becomes negatively charged. If an atom loses electrons it will become positively charged.
35
Heat
Energy that transfers from one object to another due to a difference in temperature between them. Heat always flows from the warmer to the cooler object.
36
Orbital Diagram
Shorthand method of writing the location of electrons by orbital.
36
Activation Energy
The minimum energy required to start a chemical reaction
38
Rules for Writing Lewis Dot Structures
1. Total the Valence Electrons and Update during each step.
2. Form a single bond between the central atom and each surrounding atom (each bond uses two electrons).
3. Place electrons around the outer atoms until you run out or they each have eight electrons around them. (Hydrogen and Helium can only have two electrons around them).
4. Place any remaining electrons around the central atom.
5. Does each atom other than hydrogen have eight electrons around it?
6. Rearrange (share) electrons so that all atoms other than hydrogen have eight electrons around them.
39
Which has the smallest atomic radius?
A. fluorine B. chlorine C. bromine D. iodine
F
(least number of energy levels)
39
Half Life
The time it takes for the amount of a radioactive element to decay by half.
39
immiscible
Will not dissolve in water.
40
What at problems where there with the Rutherford model of the atom?
41
In the equation:
2Na + Cl₂→ 2NaCl
How many grams of NaCl are produced from 3.75 moles of Cl₂?
3. 75 moles Cl₂ x ( 2moles NaCl/1 mole Cl₂ ) x ( 58.5g NaCl / 1 mol NaCl =
438. 75 g NaCl
41
pH Scale
pH = (-) log [H+]
pOH = - log [OH-]
pH + pOH = 14
[H+] = 2nd Log (-) pH
[OH-] = 2nd Log (-) pOH
[H+] [OH-] = 1E-14
42
Products
## Footnote
Substance(s) being made. Right side of the equation.
43
Exothermic
An exothermic reaction is a chemical reaction that releases energy in the form of light or heat.
44
Anion
Negatively charged Ions. Anions are negatively charged because they have gained an electron(s) (electrons are negative). In general, anions are nonmetals.
45
Predict the element formed and write a balanced nuclear equation for the following statement:
Uranium-238 decays by alpha emission to form \_\_\_\_\_\_\_\_\_\_.
46
Solvent
The component doing the dissolving.
The component in the largest amount.
Water is the universal solvent.
46
Endothermic
Reaction in which the system absorbs energy from its surroundings in the form of heat.
47
Heat of Fusion
The energy required to change a gram of a substance from the solid to the liquid state without changing its temperature.
48
Le Chatlier's Principle
48
## Footnote
Properties of Bases
## Footnote
Bitter and slippery
49
Kinetic Energy
The energy an object has due to its motion.
50
What is the orbital diagram for Silicon?
51
Nucleus
The center of an atom. Contains the protons and neutrons. Since neutrons have no charge and protons are positively charged, the nucleus has an overall positive charge.
51
Gas
One of four main states of matter, composed of molecules in constant random motion. Unlike a solid, a gas has no fixed shape and will take on the shape of the space available. Unlike a liquid, the intermolecular forces are very small; it has no fixed volume and will expand to fill the space available.
52
## Footnote
Vapor Pressure
## Footnote
The pressure of the vapor resulting from evaporation of a liquid (or solid) above a sample of the liquid (or solid) in a closed container.
53
How much solute will not dissolve if I have 80g of Potassium Nitrate in 100g of Water at 30C?
35g Potassium Nitrate
54
55
What is the electron configuration of Krypton?
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶
[Ar] 4s² 3d¹⁰ 4p⁶
55
## Footnote
Properties of Acids
## Footnote
Sour
56
## Footnote
What is the molarity of a solution in which 58 g of NaCl are dissolved in 1.0 L of solution?
## Footnote
1.0 M
57
Polyatomic Ionic Formulas
1. Metal written first, Polyatomic Ion is written second 2. Use criss-cross method and subscripts to insure sum of charges = 0 3. When using a subscript for polyatomic ions: place a parentheses around the polyatomic formula, put the subscript outside of the parentheses
57
Law of Conservation of Mass
Mass is neither created or destroyed during a chemical reaction - it is conserved.
Mass Reactants = Mass Products
57
How does a reaction at equilibrium shift when reactants are removed?
Reaction shifts to the left to make more reactant.
57
How does a reaction at equilibrium shift when products are removed?
Reaction shifts to the right to make more product.
58
IUPAC Rules for Alkane Nomenclature
IUPAC Rules for Alkane Nomenclature
1. Find and name the longest continuous carbon chain.
2. Identify and name groups attached to this chain.
3. Number the chain consecutively, starting at the end nearest a substituent group.
4. Designate the location of each substituent group by an appropriate number and name.
5. Assemble the name, listing groups in alphabetical order using the full name (e.g. cyclopropyl before isobutyl).
The prefixes di, tri, tetra etc., used to designate several groups of the same kind, are not considered when alphabetizing.
59
Which has the largest atomic radius?
K
Na
Rb
Li
Rb
| (most energy levels)
59
Plastic
Long chain organic compounds that can be molded or formed into shapes, films, or fibers.
61
What do the symbols in the following equation represent:
Q=mC(Delta T)
63
How many atoms are in 2.7 moles of iron?
2.7 moles Fe x (6.02x10²³ atoms Fe/1 mole Fe) =
## Footnote
1.6 x 10²⁴atoms Fe
63
## Footnote
Diatomic Molecule
## Footnote
If these elements appear by themselves in an equation, the must be in pairs with the subscript 2.
63
## Footnote
Na + MgCl₂ →
## Footnote
2Na + MgCl₂ →2NaCl + Mg Single Replacement
63
Intramolecular Forces
Attractive forces WITHIN a molecule (Ionic & Covalent Bonds)
These bonds are stronger than intermolecular forces.
64
miscible
Will dissolve in water.
65
Predict the products of and write a balanced equation from the following statement:
HCl + Zn →
2HCl + Zn →ZnCl₂ + H₂
Single Replacement
66
Principle Energy Level
The possible locations around an atom where electrons having specific energy values (quantum number) may be found. Divided into sublevels s, p, d, and f.
67
## Footnote
K + Cl₂ →
## Footnote
2K + Cl₂ →2KCl Synthesis
68
Weak Acid
Weak acids are not fully ionized or dissociated.
69
Solute
The component being dissolved.
The component in the smallest amount.
70
How can hydrogen bonding affect the properties of polymers?
Different side groups on the polymer can lend the polymer to ionic bonding or hydrogen bonding between its own chains. These stronger forces typically result in higher tensile strength and higher crystalline melting points.
72
Radioactivity
When the nucleus of an atom becomes unstable and decays. This tends to happen to large atoms (larger than atomic number 83) becuase the number of protons to neutrons in the nucleus is unbalanced.
73
Boiling
Boiling is the rapid vaporization of a liquid, which occurs when a liquid is heated to its boiling point.
74
Predict the products of and write a balanced equation from the following statement:
CuCl₂ + H₂S →
CuCl₂ + H₂S →CuS + 2HCl
Double Replacement
74
Balance the following chemical equation:
\_\_C + \_\_O₂ → \_\_CO₂
## Footnote
1C + 1O₂ → 1CO₂
74
## Footnote
Mg + I₂ →
## Footnote
Mg + I₂ →MgI₂ Synthesis
75
IUPAC Rules for Alkyne Nomenclature
IUPAC Rules for Alkyne Nomenclature
1. The yne suffix (ending) indicates an alkyne or cycloalkyne.
2. The longest chain chosen for the root name must include both carbon atoms of the triple bond.
3. The root chain must be numbered from the end nearest a triple bond carbon atom. If the triple bond is in the center of the chain, the nearest substituent rule is used to determine the end where numbering starts.
4. The smaller of the two numbers designating the carbon atoms of the triple bond is used as the triple bond locator.
5. If several multiple bonds are present, each must be assigned a locator number. Double bonds precede triple bonds in the IUPAC name, but the chain is numbered from the end nearest a multiple bond, regardless of its nature.
76
Roman Numerals
Above groups/families Roman Numerals represent the number of valence electrons.
When writing the name of an ionic compound which includes a transition metal a Roman Numeral is used to denote the quantity of positive charge associated with that transition metal.
78
What element is represented by the following orbital diagram?
Fluorine
78
Polymers
Very long chains of repeating units (monomers)
| (Simplest polymer: -(-CH₂CH₂CH₂CH₂CH₂-)- =Polyethylene)
79
## Footnote
Bases
Electron donator.
Yeilds OH- in solution.
H+ Ion acceptor.
79
Heat Capacity
Heat capacity is the amount of heat energy required to change the temperature of a substance
80
Atmospheric Pressure
Pressure resulting from the collisions of atoms and molecules with objects.
81
How does a reaction at equilibrium shift when more reactants are added?
Reaction shifts to the right to make more product.
82
When is a photon emitted from an atom?
A photon is emitted as an excited electron returns to its ground state.
83
9.08 g of Al₂0₃ is equal to how many moles?
9.08 g Al₂0₃ x (1 mole Al₂0₃/101.96 g Al₂0₃) = 0.0890 moles Al₂0₃
84
Properties of Ionic Compounds
Crystalline Structure High Melting Point Rigid Strong Bond Conduct Electricity (when dissolved or melted) Good Insulators
85
9.08 g of Al₂0₃ is equal to how many moles?
9.08 g Al₂0₃ x (1 mole Al₂0₃/101.96 g Al₂0₃) = 0.0890 moles Al₂0₃
86
Which of the following elements has the largest atomic radius?
A. nitrogen B. oxygen C. fluorine D. neon
N
(least number of protons attracting the electrons on the energy levels)
87
Ionic Compound Formulas
1. Metal written first Nonmetal written second 2. Use criss-cross method and subscripts to insure sum of charges = 0
87
Liquid
Liquid is one of the four fundamental states of matter, and is the only state with a definite volume but no fixed shape. A liquid is made up of tiny vibrating particles of matter, such as atoms, held together by intermolecular bonds.
88
Summarize how to write formulas for ionics and covalents.
89
Gases deviate most from ideal behavior under conditions of very ______ temperature and very _______ pressure
low; high
90
Describe the Bohr Model of an atom.
92
VSPER
Valence Shell Electron Pair Repulsion
Most important aspect of determining the shape or geometry of a molecule. The molecule will adopt the shape which minimizes te electron pair repulsion.
92
Chemical Potential Energy
The energy stored in chemical bonds. (Lines with no Slope on the Heating Curve)
94
What is the IUPAC name of
94
How many atoms are in 2.7 moles of iron?
2.7 moles Fe x (6.02x10²³ atoms Fe/1 mole Fe) =
## Footnote
1.6 x 10²⁴atoms Fe
96
Kinetic Energy
In Thermochemistry, the energy an object possesses due to its change in temperature is called kinetic energy. (Sloped lines on the Heating Curve)
97
## Footnote
How many grams of KNO₃ should be used to prepare 2.00 L of a .500 M solution?
## Footnote
101 g KNO₃
98
How many moles are in 3.67 x 10²⁴ molecules of SO₂?
3. 67 x 10²⁴ molecules SO₂ x (1 mole SO₂/6.02x10²³ molecules SO₂)=
6. 10 moles SO₂
100
Compound Formation
Compounds are the result of the formation of chemical bonds between two or more different elements, whose atoms lose, gain or share valence electrons to complete their outer shell and attain a noble gas configuration.
This tendency of atoms to have eight electrons in their outer shell is known as the octet rule.
101
What is the effect of pressure on solubility?
102
## Footnote
Directly Proportional
## Footnote
As the value of one variable increases, the value of the other increases and vice-versa (e.g., temperature and volume, temperature and pressure)
103
Period
Horizontal row on the periodic table.
104
What is the orbital diagram for Sulfur?
105
Stereo Isomers
A carbon that has 4 different things attached to it is considered "chiral"
There is a mirror image of the molecule, they are stereoisomers or enantiomers, bend light oppositely (optically active)
106
Definition of a balanced chemical equation:
no atoms are lost or gained
the number of reacting atoms is equal to the number of product atoms
108
Polar Covalent
A type of covalent bond between two atoms in which electrons are shared unequally. Because of this, one end of the molecule has a slightly negative charge and the other a slightly positive charge.
110
Empirical Formula
Empirical Formula - A formula that gives the simplest whole-number ratio of atoms in a compound.
111
How to speed up the Rate of a Reaction
1. Shrink the container or increase concentration. 2. Increase the number of particles. 3. Speed up the particles by adding heat. 4. Break up clumps into individual particles. 5. Use a catalyst.
112
Balance the following chemical equation:
\_\_C₂H₆ + \_\_O₂ → \_\_CO₂ + \_\_H₂O
## Footnote
2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O
112
Calculate the concentration of the solute in the diluted solution: 1.00mL of 0.50M NH4Cl diluted to 250.0mL
M1V1 = M2V2
M1 = (M2V2)/(V1)
M1 = (0.5 M x 1 ml)/(250 ml)
M1 = 0.0020M
113
Calorimeter
The device used to measure the quantity of heat flow in a reaction.
115
A compound contains 21.6 g of silver and 3.21 g of sulfur, what is the percent composition?
21.6/(21.6+3.21) = 87.1% Ag
## Footnote
3.21/(21.6+3.21) = 12.9%S
116
## Footnote
What is the molarity of a solution in which 10.0 g of AgNO₃ is dissolved in 500 ml of solution?
## Footnote
0.118 M AgNO₃
117
How much more solute could I dissolve if I have 10g of Potassium Nitrate in 100g of Water at 30C?
~32g of Potassium Nitrate
118
Bond Pairs
Pairs of electrons which form a bond.
120
Electrons
Negatively charged subatomic particles. Electrons can be found in the space around the nucleus. This area is often called the electron cloud. Electrons have NO mass.
122
Electronegativity
Tendency of an atom to attract electrons to itself when it combines with another element.
Electronegativity increases from left to right across a period on the periodic table.
Electronegativiy decreases as you go down a group on the periodic table.
123
Double Displacement
There are two compounds on each side of the equation. During a DD Reaction the positively charged cations in each compound switch places. Also known as Precipitate Reaction. For the reaction to occur a precipitate must be formed. The solubility rules or table are used to confirm this.
AB + CD --\> AD + CB
123
## Footnote
How would you calculate the pH of a solution with a hydronium concentration of 1E-5 M?
## Footnote
pH = -log 1E-5 =5
124
How much solute will not dissolve if I have 80g of Potassium Nitrate in 100g of Water at 30C?
35g Potassium Nitrate
125
Boiling Point
The temperature at which the vapor pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding environmental pressure. At the boiling point the gas and liquid phase exist in equilibrium.
128
Proton
Positively charged subatomic particle. Contained inside the nucleus. Neutrons and Protons weigh roughly the same (1 atomic mass unit).
129
Standard Temperature & Pressure (STP) - a shorthand method that can be used when writing word problems.
130
Potential Energy
Potential energy is that energy which an object has because of its position. It is called potential energy because it has the potential to be converted into other forms of energy, such as kinetic energy. (Lines with no Slope on the Heating Curve
131
Naming Covalent Compunds
You DO NOT use the “criss-cross” method. Use prefixes. The only time you do not use a prefix is when there is only one of the first nonmetal. Remember : 1-Mono 2-Di etc...
131
What happens to an exothermic reaction when temperature is lowered?
Heat is a product, so the reaction shifts to the right to make more product.
132
Dilute
Relatively small amount of solute.
134
Percent Composition
The amount of a given element in a compound.
135
Subscript
Small numbers to the lower right of chemical symbols. Represent the number of atoms of each element in a molecule.
O₂
137
Empirical Formula
Empirical Formula - A formula that gives the simplest whole-number ratio of atoms in a compound.
138
Melting Point
The melting point of a solid is the temperature at which it changes state from solid to liquid at atmospheric pressure. At the melting point the solid and liquid phase exist in equilibrium.
139
Catalyst
A substance that causes a chemical reaction to happen more quickly.
141
Polyatomic Ion
Polyatomic Ions ions made up of more than one atom; acts as an individual ion in a compound; charge applies to entire group of atoms.
143
Isotopes
Atoms of the same element which have a different Mass Number due to a differing number of neutrons.
The symbol for an Isotope is the element symbol followed by the Mass Number, for example:
Na-23 and Na-24
144
In the equation:
NH₄NO₃ → N₂O + 2H₂O
How many grams of H₂O are produced if you are given 50 g of NH₄NO₃?
50 g NH₄NO₃ x (1 mole NH₄NO₃ /80.04 gNH₄NO₃) x
(2 moles H₂O / 1 mole NH₄NO₃) x (18.02 g H₂O/ 1 mol H₂O) =
22.5 g H₂O
144
Freezing
Freezing or solidification is a phase transition in which a liquid turns into a solid when its temperature is lowered below its freezing point. (Liquid to Solid)
145
What is the electron configuration of Beryllium?
1s² 2s²
145
What are the four Quantum Numbers? What do they stand for?
145
Which has the highest electronegativity?
Na
Al
S
Cl
Cl
145
What is the IUPAC name of
145
Condensation (Precipitation)
The conversion of a vapor or gas to a liquid.
146
Percent Mass
147
Solubility Curve
Used to determine the mass of solute in 100g (100ml) of water at a given temperature.
149
What is wrong with the following orbital diagram?
This diagram does not follow Hund's Rule. Electrons must fill energy levels from lowest to highest.
150
Write a chemical equation from the following word equation:
Water decomposes into hydrogen gas and oxygen gas:
## Footnote
2H₂O → 2H₂ + O₂
151
## Footnote
What is the Lewis dot structure of Br₂?
151
How to PREDICT PRODUCTS of the 4 Main Types of Chemical Reactions.
153
Precipitate
## Footnote
Formation of a solid in a solution.
154
The Mole
A unit of measure.
1 mole = the atomic mass of an element
and
1 mole = 6.02 E23 atoms, particles, molecules, units, etc.
155
Write a chemical equation from the following word equation:
Sodium reacts with fluorine to form sodium fluoride
## Footnote
2Na + F₂--\> 2NaF
156
What is the IUPAC name of
157
## Footnote
pH, pOH, [H+], [OH-] Calculation Road Map
158
Write a chemical equation from the following word equation:
Magnesium reacts with aluminum chloride to form magnesium chloride and aluminum.
## Footnote
3Mg + 2AlCl₃--\> 3MgCl₂ + 2Al
160
Atomic Size
Size of an atom.
Depends on the number of energy levels and the overall charge of the nucleus (the nucleus's charge depends on the number of protons).
As you go down a group you are adding energy levels (rings around the nucleus) so the atomic size increases.
As you go across a period the atomic size decreases because the energy levels (which contain negative electrons) are increasingly attracted to the positive protons in the nucleus of the atom.
162
## Footnote
Coefficients
## Footnote
The large numbers in front of chemical formulas. Represent the number of molecules of the substance in the reaction.
2Cl
163
What is the solubility of Potassium Nitrate at 50 C?
86 g Potassium Nitrate
167
Diatomic Molecules
If these elements appear by themselves in an equation, the must be in pairs with the subscript 2.
(H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂)
170
Synthesis or Combination
Two or more elements combine to make one product.
A + B --\> AB
171
## Footnote
Calculate the pH of a solution with a Hydronium ion concentration of 1.3E-9 M
## Footnote
pH= -log(1.3E-9) =8.89
172
Covalent/Molecular Bonding
Nonmetals SHARE electrons
173
Cation Formation
Metals give electrons to become like the nearest noble gas.
173
Percent Composition
The amount of a given element in a compound.
173
Vaporization (Evaporation)
The conversion of a liquid into a gas.
174
Actual Yeild
The amount of product which is actually formed when the reaction is carried out in the laboratory.
175
Theoretical Yeild
The maximum amount of product that is CALCULATED to be formed from the given amounts of reactants.
177
Dimensional Analysis
A method of mathmatical analysis which involves converting between units.
178
## Footnote
Inversely Proportional
## Footnote
As the value of one variable increases, the value of the other decreases and vice-versa (e.g., pressure and volume)
179
What is the electron configuration of Zinc?
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰
[Ar] 4s² 3d¹⁰
181
Unsaturated Hydrocarbons
Have double or triple bonds (they can react and get more atoms in the molecule; alkenes and alkynes)
182
Steps for finding Empirical Formula from percent composition
184
How many moles are in 3.67 x 10²⁴ molecules of SO₂?
3. 67 x 10²⁴ molecules SO₂ x (1 mole SO₂/6.02x10²³ molecules SO₂)=
6. 10 moles SO₂
185
Molarity
Expression of Concentration ( [] )
186
Kinetic Molecular Theory of Gases
1. Gases consist of tiny particles.
2. These particles are so small, their volume can be assumed to be negligible (zero).
3. The particles are in constant, random motion, colliding with the walls of the container. These collisions with the walls cause the pressure exerted by the gas.
4. The particles are assumed not to attract or repel each other.
5. The kinetic energy of the particles is directly proportional to the Kelvin temperature of the gas.
(As the temperature increases the energy increases)
187
Octet Rule
All atoms other than hydrogen and helium bond to fill their valence shell with eight electrons.
189
Fission
191
Symbols used in writing chemical equations
192
Parent Chain
Longest carbon chain
193
What happens to an endothermic reaction when temperature is lowered?
Heat is a reactant, so the reaction will shift to the left to make more reactant.
195
## Footnote
NEUTRALIZATION REACTION
## Footnote
When an acid and base react in solution to form salt and water.
197
How would removing ammonia, NH3, affect the equilibrium of the following reaction?
N2 + 3H2 ⇔ 2NH3 + heat
198
Concentrated
A relatively large amount of solute.
200
1. What is the IUPAC name for the following compound?
2,2,4-trimethylpentane
202
How would temperature and pressure affect the equilibrium of the following reaction?
N2 + 3H2 ⇔ 2NH3 + heat
203
Colligative Properties
properties that depend on the concentration of solute, but not on the chemical identity of the solute.
(Vapor Pressure Lowering, Boiling Point Elevation, Freezing Point Depression, and Osmotic Pressure)
205
Covalent or Molecular Compounds (Molecules)
The first word gives name of the element that appears first preceded by a prefix that shows number of atoms in that element. Do NOT use the prefix mono before the name of the first element. The second word consists of a prefix designating the number of atoms of that element, the stem name of the second element, and the suffix -ide.
206
Anion Formation
Nonmetals recieve electrons when forming an ionic compound so that their electron configuration will be like the nearest noble gas.
208
## Footnote
When adding and subtracting with Significant Figures, \_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_.
## Footnote
keep the least number of decimal places
209
A compound contains 21.6 g of silver and 3.21 g of sulfur, what is the percent composition?
21.6/(21.6+3.21) = 87.1% Ag
## Footnote
3.21/(21.6+3.21) = 12.9%S
210
How much solute would crystallize out of solution once a solution of 85g of Potassium Nitrate in 100g of Water at 50C is cooled to 20C?
55g Potassium Nitrate
211
Monomer
the basic unit of the polymer
211
What is a condensation (step-growth) polymer?
A polymer formed from two different monomers that are joined together with the elimination of a small molecule. Nylon is a condensation polymer.
212
Anions
Anions are named by adding the suffix -ide to the stem of the name of the nonmetal from which they are derived. N³⁻ = Nitride O²⁻ = Oxide F⁻ = Flouride
213
Strong Acid
Completely dissociates.
214
Predict the products of and write a balanced equation from the following statement:
Mg + I₂ →
Mg + I₂ →MgI₂
Synthesis
215
Calorimetry Steps
1. Use a well known substance (water) 2. Run the reaction and let heat go into or out of the water 3.Measure the waters change in temperature
217
How does a reaction at equilibrium shift when more products are added?
Reaction shifts to the left to make more reactant.
218
Which has the largest atomic radius?
K
Na
Rb
Li
Rb
| (most energy levels)
219
Temperature
A measure of the intensity of heat, i.e. the hotness or coldness of a sample. or object.
221
What are both of the symbols for each of the following types of nuclear decay:
Alpha, Beta, Gamma
222
Hydrogen
Group 1 Nonmetal
223
Which has the highest electronegativity?
Na
Al
S
Cl
Cl
225
Decomposition Reaction
One reactant breaks down into two or more elements.
AB --\> A + B
226
Phase Changes
227
## Footnote
Subscripts
## Footnote
Small numbers to the lower right of chemical symbols. Represent the number of atoms of each element in a molecule.
0₂
228
Cations
Positively Charged Ions. In general, cations are metals.
230
Thermoplastics
Long chains of organic compounds which have NO cross-links. These may be melted and recycled into new shapes (products).
233
Thermosetting Plastics
Long Chain Organic Compounds which HAVE cross-links which inhibits melting and thus the recycling of these products. Once set in a certain shape they can not be remolded.
(or thermosetting, plastics are synthetic
materials that strengthen during being
heated, but cannot be successfully
remolded or reheated after their initial
heat-forming. )
234
Homologous
A Homologous Series is a group of organic chemical compounds, usually listed in order of increasing size,
that have a similar structure (and hence also similar properties) and whose structures differ only by the number of CH2 units in the main carbon chain.
(The simplest example of a homologous series in organic chemistry is that of alkanes)
235
Which has the smallest atomic radius?
A. fluorine B. chlorine C. bromine D. iodine
F
(least number of energy levels)
236
What is the effect of temperature on solubility?
237
## Footnote
Hydronium Ion
## Footnote
[H+]
239
Hund's Rule
When filling sublevels other than s, electrons are placed in individual orbitals before they are paired up.
Electrons fill like people do on a bus. You would never sit right next to someone you do not know if there are free seats available, unless of course all the seats are taken then you must pair up.
So, when working with the p sublevel, electrons fill like this...up, up, up....down, down, down...
241
Balance the following chemical equation:
\_\_Al₂(SO₄)₃ + \_\_Ca(OH)₂ → \_\_Al(OH)₃ + \_\_CaSO₄
## Footnote
Al₂(SO₄)₃ + 3Ca(OH)₂ → 2Al(OH)₃ + 3CaSO₄
243
## Footnote
What would the Lewis dot structure of OC be?
244
Collision Theory
Molecules have to collide in a certain way with a certain amount of activation energy in order to form a new product.
245
Solidification
Freezing or solidification is a phase transition in which a liquid turns into a solid when its temperature is lowered below its freezing point. (Liquid to Solid)
247
Isomers
Same formula, different structure
249
Which has the lowest electronegativity?
F
I
Br
Cl
I
251
## Footnote
NaOH + HClO₄ →
## Footnote
NaOH + HClO₄ →NaClO₄+ H₂O Double Replacement
254
What is the molarity of a solution made by dissolving 2.5g of NaCl in enough water to make 125 ml of solution?
255
The heat lost in a reaction (Q lost) = ?
Q lost = Q gained
256
Valence Electrons
Electrons found on the outer energy level. Represented by Roman Numerals. Transition Elements do not have Roman Numerals which denote their valence electrons.
258
"Like Dissolves Like"
Not everything will dissolve in everything. Polar solvents dissolve polar solutes. Nonpolar solutes dissolve nonpolar solutes. Water is polar so it dissolves polar solutes.
259
Sublevel
Principal energy levels are broken down into sublevels designated s, p, d, or f upon which electrons travel.
261
In the equation:
2Na + Cl₂→ 2NaCl
How many moles of Na are needed if 285 g of NaCl are produced?
285g NaCl x (1 mole NaCl/ 58.5 g NaCl) x (2 moles Na / 2 moles NaCl) =
4.87 moles Na.
262
What is the electron configuration of Calcium?
1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
[Ar] 4s²
263
## Footnote
Diatomic Molecule
## Footnote
If these elements appear by themselves in an equation, the must be in pairs with the subscript 2.
264
Limiting Reagant/Reactant
Determines the amount of product that can be formed.
The reactant/reagant that is completely used up during a reaction.
(In the attached picture, beef patties are limiting how many hamburgers we can make.)
265
IUPAC Rules for Alkene Nomenclature
1. The ene suffix (ending) indicates an alkene or cycloalkene.
2. The longest chain chosen for the root name must include both carbon atoms of the double bond.
3. The root chain must be numbered from the end nearest a double bond carbon atom. If the double bond is in the center of the chain, the nearest substituent rule is used to determine the end where numbering starts.
4. The smaller of the two numbers designating the carbon atoms of the double bond is used as the double bond locator. If more than one double bond is present the compound is named as a diene, triene or equivalent prefix indicating the number of double bonds, and each double bond is assigned a locator number.
266
Dimensional Analysis
A method of mathmatical analysis which involves converting between units.
267
Neutrons
Neutrally charged subatomic particle. (No charge) Contained inside the nucleus. Neutrons and Protons weigh roughly the same (1 atomic mass unit)
269
Intermolecular Forces
Attractive forces BETWEEN molecules.
Van Der Waals or London Dispersion Forces are the weakest type of intermolecular force and hydrogen bonds are the strongest.
270
Excess reagent/reactant
The reactant/reagant which is not completely used up during a reaction.
271
Predict the products of and write a balanced equation from the following word statement:
Methane (C3H8) reacts with Oxygen gas to form \_\_\_\_\_\_\_\_\_\_.
272
Positive Exponents in Scientific Notation
Postive Exponents mean move the decimal to the right when changing a number to standard notation.
273
## Footnote
HCl + Zn →
## Footnote
2HCl + Zn →ZnCl₂ + H₂ Single Replacement
274
What is the orbital diagram for Nitrogen?
275
Electronegativity
Tendency of an atom to attract electrons to itself when it combines with another element.
Electronegativity increases from left to right across a period on the periodic table.
Electronegativity decreases as you go down a group on the periodic table.
276
Ideal Gas Law
278
## Footnote
Coefficients
## Footnote
The large numbers in front of chemical formulas. Represent the number of molecules of the substance in the reaction.
2Cl
279
Which has the lowest electronegativity?
F
I
Br
Cl
I
280
If the [H+] =1E-5 in a Coca-Cola what is the concentration of [OH-]? Is the solution acidic, basic, or neutral?
[H+] [OH-] = 1E-14
[OH-] = (1E-14) / ([H+])
[OH-] = (1E-14) / (1E-5)
[OH-] = 1E-9
The solution is acidic because 1E-5 is less than 1E-7.
281
Is the solution unsaturated, saturated, or supersaturated, if I have 30 g of Potassium Nitrate in 100 g of Water at 30 C?
Unsaturated
282
What is the IUPAC name of
283
What are the conversions between pressure units? (What amount of atm equals what amount of kPa, psi, torr, etc)
1 atm = 760 torr = 760 mmHg = 101.3 kPa
284
What is the percent composition of oxygen in CuSO₄?
mass of CuSO₄= 63.546 + 32.065 + 4(15.999) = 159.7 grams
## Footnote
Mass of oxygen = 4(15.999) = 64 grams
(64/159.7) x (100%) = 40.1. %O
286
Summarize the rules for naming ionic and covalent compounds.
288
Fusion
289
Is the solution unsaturated, saturated, or supersaturated, if I have 30 g of Potassium Nitrate in 100 g of Water at 30 C?
Unsaturated
290
Saturated
Contains as much solute as it can at that temperature.
If you add more solute it will not dissolve.
292
Synthetic or Engineered Polymers
1. Thermoplastics
2. Thermosetting plastics
3. Elastomers (Rubber)
293
What is the orbital diagram for Oxygen?
294
What mass of copper (II) sulfide is equal to 4.79 x 10²² atoms?
4. 79 x 10²² atoms CuS x (1 mole CuS/6.02x10²³ atoms CuS) x (95.61 g CuS/1 mole CuS) =
7. 61 g CuS
295
Heat
A form of energy that flows between two samples of matter because of their differences in temperature.
296
Supersaturated
Contains more solute than it should at that temperature. Ussually the solution is heated to allow that much solute to dissolve and then cooled back down.
If you add solute to a supersaturated solution the "super" part (amount over the saturation point) will crystallize out of solution.
297
Polymerization
the process by which the synthesis of polymers occurs.
(the process by which monomers build polymers)
298
Heating Curve
A plot of the temperature versus time.
299
Saturated Hydrocarbons
300
## Footnote
What is the Lewis dot structure of XeF₄?
301
## Footnote
CuCl₂ + H₂S →
## Footnote
CuCl₂ + H₂S →CuS + 2HCl Double Replacement
302
alkane
hydrocarbon with single bonds between all carbons
303
Predict the products of and write a balanced equation from the following statement:
C₄H₆ + O₂ →
2C₄H₆ + 7O₂ →4CO₂ + 6H₂O
Combustion
304
Pauli Exclusion Principle
If there are two electrons in an orbital, they must have opposite (paired) spins.
305
Determine the conjugate acid/base pairs.
NH3 + H2O ---\> NH4+ + OH-
306
Solution
Solute + Solvent
307
Given the following reaction: S₈ + 4Cl₂ → 4S₂Cl₂. If there is 200g of sulfur and 100g of chlorine, what mass of disulfur dichloride will be produced?
1. Perform a mass-to-mass calculation between sulfur and disulfur dichloride.
200g S₈ x (1mol S₈ / 256.5g S₈) x (4mol S₂Cl₂ / 1 mol S₈) x (135g S₂Cl₂ / 1 mol S₂Cl₂) = 421g S₂Cl₂.
2. Perform a mass-to-mass calculation between chlorine and disulfur dichloride.
100g Cl₂ x (1mol Cl₂ / 70.91g Cl₂) x (4mol S₂Cl₂ / 4 mol Cl₂) x (135g S₂Cl₂ / 1 mol S₂Cl₂) = 190.4g S₂Cl₂.
The limiting reactant is chlorine (Cl₂ ) because... it produced only 190.4g S₂Cl₂.
The excess reactant is sulfur (S₈) because... it would have produced 421g S₂Cl₂.
308
What mass of copper (II) sulfide is equal to 4.79 x 10²² atoms?
4. 79 x 10²² atoms CuS x (1 mole CuS/6.02x10²³ atoms CuS) x (95.61 g CuS/1 mole CuS) =
7. 61 g CuS
309
In the equation:
C₃H₈ + 5O₂→ 3CO₂+ 4H₂O
How many moles of CO₂are produced when 7 moles of C₃H₈ are burned?
7 moles C₃H₈ x (3 moles CO₂ / 1 mole C₃H₈) =
21 moles CO₂
310
Steps for finding Empirical Formula from percent composition
311
What is the percent composition of oxygen in CuSO₄?
mass of CuSO₄= 63.546 + 32.065 + 4(15.999) = 159.7 grams
## Footnote
Mass of oxygen = 4(15.999) = 64 grams
(64/159.7) x (100%) = 40.1. %O
312
Evaporation versus Boiling
313
thermochemical equation
thermochemical equation
- balanced chemical equations with heat changes taken into account
- think of it as (and treat it as) a reactant or a product
314
Enthalpy (H)
-the heat content of a system (or chemical reaction)
315
CaO (s) + H₂O (l) → Ca(OH)₂(s) + 65.2 kJ
Change in H = - 65.2 kJ
Endo or Exo?
Exo
heat is a product and change in H is negative
316
2 NaHCO₃ (s) + 129 kJ → Na₂O (s) + H₂O (g) + CO₂(g)
Change in H = +129 kJ
Endo or Exo ? How to tell?
Endo
heat is a reactant and change in H is positive
317
Endo or Exothermic?
## Footnote
Exothermic reaction graph
318
five ways to tell if a reaction is exothermic
graph slopes down
heat is given off (to the surroundings)
change in H is negative
temperature of "surroundings" increases
heat is written as a product
319
five ways tell if a reaction is endothermic
graph slopes up
heat is absorbed (from the surroundings)
change in H is positive
temperature of "surroundings" decreases
heat is written as a reactant
320
Phase Diagram most substances (e.g. CO2)
slight positive slope; more dens in (l) phase;
triple point: s, l, g in equilib.
critical point: T and P above which there is no distinction between l and g
321
## Footnote
Phase Diagram H2O
negative slope bc H2O has lower density in (s) than (l); increase in P favors more dense version.
322
Boyle's Law
The pressure and volume of a gas are inversely proportional at a constant temperature.
P1V1 = P2V2
323
Charles' Law
The volume of a gas is directly proportional with the KELVIN temperature at a constant pressure.
V1T2 =V2T1
324
Kinetic Molecular Theory of Gases
1. Gases consist of tiny particles.
2. These particles are so small, their volume can be assumed to be negligible (zero).
3. The particles are in constant, random motion, colliding with the walls of the container. These collisions with the walls cause the pressure exerted by the gas.
4. The particles are assumed not to attract or repel each other.
5. The kinetic energy of the particles is directly proportional to the Kelvin temperature of the gas.
(As the temperature increases the energy increases)
325
Standard Temperature & Pressure (STP) - a shorthand method that can be used when writing word problems.
326
Barometer
Measures atomospheric pressure.
327
Combined Gas Law
P1V1T2 = P2V2T1
328
Ideal Gas Law
329
Dalton's Law of Partial Pressures
330
Molar Volume of a Gas
1 mole of any gas at STP = 22.4 Liters
331
Kinetic Energy
The energy an object has due to its motion.
332
Atmospheric Pressure
Pressure resulting from the collisions of atoms and molecules with objects.
333
What are the conversions between pressure units? (What amount of atm equals what amount of kPa, psi, torr, etc)
1 atm = 760 torr = 760 mmHg = 101.3 kPa
334
If 2.5g of sulfur hexafluoride is introduced to an evacuated 500.0mL container at 83\*C, what is the pressure, in atmospheres, inside the container?
PV = nRT
2.5g SF6 (1mole SF6/ 146.08 g SF6) = 0.017 moles SF6
P = (nRT)/V
P =[ (0.017 moles)x(0.08205746 L atm/K mol)x(83+273K)/(.5 L)
P =0.99atm
335
If the total pressure of a two gas system is 100 torr and the partial pressure of one gas is 70 torr, what is the pressure of the other gas?
Pother = 30 torr
336
Gases **_deviate_** most from ideal behavior under conditions of very ______ temperature and very _______ pressure
low; high
337
According to _____________ law, pressure and volume are ______________ proportional provided all other factors remain constant. Mathematically, this means that their ____________ is a constant.
Boyle's; inversely; temperature
338
If pressure is constant, the volume of a sample of gas _____________ as temperature increases.
increases
339
At constant pressure, the volume of a sample of gas is ___________ proportional to temperature as measured on the ____________ temperature scale
directly; Kelvin
340
A balloon with a volume of 2.0 L is filled with a gas at 3 atmospheres. If the pressure is reduced to 0.5 atmospheres without a change in temperature, what would be the volume of the balloon?
Since the temperature does not change, Boyle's law can be used. Boyle's gas law can be expressed as:
## Footnote
PiVi = PfVf (i=initial and f=final)
To find the final volume, solve the equation for Vf:
Vf = PiVi/Pf
Vi = 2.0 L
Pi = 3 atm
Pf = 0.5 atm
Vf = (2.0 L)(3 atm)/(0.5 atm)
Vf = 6 L/0.5
Vf = 12 L
Answer:
The volume of the balloon will expand to 12 L.
341
A 600 mL sample of nitrogen is heated from 27 °C to 77 °C at constant pressure. What is the final volume?
The first step to solving gas law problems should be converting all temperatures to absolute temperatures. This is the most common place mistakes are made in this type of homework problem.
## Footnote
K = 273 + °C
Ti = initial temperature = 27 °C
K = 273 + 27
Ki = 300 K
Tf = final temperature = 77 °C
K = 273 + 77
Kf = 350 K
The next step is to use Charles' law to find the final volume. Charles' law is expressed as:
ViTf = VfTi
where
Vi and Ti is the initial volume and temperature
Vf and Tf is the final volume and temperature
Solve the equation for Vf:
Vf = ViTf/Ti
Enter the known values and solve for Vf.
Vf = (600 mL)(350 K)/(300 K)
Vf = 700 mL
Answer:
The final volume after heating will be 700 mL.
342
The pressure of a mixture of nitrogen, carbon dioxide, and oxygen is 150 kPa. What is the partial pressure of oxygen if the partial pressures of the nitrogen and carbon dioxide are 100 kPA and 24 kPa, respectively?
P = Pnitrogen + Pcarbon dioxide + Poxygen
150 kPa = 100 kPa + 24 kPa + Poxygen
Poxygen = 150 kPa - 100 kPa - 24 kPa
Poxygen = 26 kPa
343
A cylinder contain a gas of volume 30 L, at a pressure of 110 kPa and a temperature of 420 K. Find the temperature of the gas which has a volume 40 L at a pressure of 120 kPa.
Vi = 30 L, Pi = 110 kPa, Ti = 420 K, Vf = 40 L, Pf = 120 kPa
Final Temperature(Tf) = PfVfTi / PiVi
Tf= (120kPa x 40L x 420K) / (110kPa x 30L)
Tf = 610.91 K
344
A cylinder contain a gas of volume 10 L, at a pressure of 80 kPa and a temperature of 200 K. Find the temperature of the gas which has a volume 20 L at a temperature of 220 K.
Vi = 10 L, Pi = 80 kPa, Ti = 200 K, Vf = 20 L, Tf = 220 kPa
Final Pressure(Pf) = PiViTf / TiVf
Pf= (80 x 10 x 220) / (200 x 20)
Pf = 44 kPa
345
## Footnote
Inversely Proportional
## Footnote
As the value of one variable increases, the value of the other decreases and vice-versa (e.g., pressure and volume)
346
## Footnote
Directly Proportional
## Footnote
As the value of one variable increases, the value of the other increases and vice-versa (e.g., temperature and volume, temperature and pressure)
347
## Footnote
Vapor Pressure
## Footnote
The pressure of the vapor resulting from evaporation of a liquid (or solid) above a sample of the liquid (or solid) in a closed container.
348
## Footnote
Ideal Gas
## Footnote
A gas that behaves exactly as the Kinetic Molecular Theory describes
Gases **_follow_** ideal behavior under conditions of very high temperatures and very low pressures
349
## Footnote
Manometer
## Footnote
Measures the pressure of a gas in a closed container.
350
## Footnote
A balloon contains 30.0 L of helium gas at 103kPa. What is the volume of the helium when the balloon rises to an altitude where the pressure is only 25.0 kPa?
## Footnote
124 L
351
## Footnote
A balloon inflated in a room at 24 c has a volume of 4.00 L. The balloon is then heated to a temperature of 58.0 c . What is the new volume if the pressure remains constant?
## Footnote
4.46 L
352
## Footnote
At 34.0 c the pressure inside a nitrogen filled tennis ball with a volume of 0.148 L is 212 kPa. How many moles of nitrogen gas are in the tennis ball?
## Footnote
0.0122 moles
