Final Exam Flashcards
(282 cards)
1
Q
Chemistry
A
The study of the properties and behavior of matter
2
Q
Matter
A
The physical material of the universe; anything that has mass and occupies space
3
Q
Property
A
Any characteristic that allows us to recognize a particular type of matter and to distinguish it from other types
4
Q
Element
A
A substance that cannot be decomposed into simpler substances. Each is composed of a unique kind (one type) of atom
5
Q
Atoms
A
Small building blocks of matter
6
Q
What are properties of matter related to?
A
The composition (kinds of atoms the matter contains) and the structure (arrangement of the atoms)
7
Q
Molecules
A
Two or more atoms joined in specific shapes
8
Q
States of matter
A
Solid, liquid, gas
9
Q
Pure substance (substance)
A
Matter that has distinct properties and a composition that does not vary from sample to sample. All substances are either elements or compounds. Has a fixed composition
Ex: water and sodium chloride
10
Q
Compounds
A
Substances composed of two or more elements; contain two or more kinds of atoms
11
Q
Mixture
A
Combinations of two or more substances in which each substance retains its chemical identity
12
Q
Law of Constant Composition/Law of Definite Proportions
A
The elemental composition of a compound is always the same. Compounds composed in the lab and the corresponding compound found in nature are the same. A pure compound has the same composition and properties under the same conditions
13
Q
Mixtures (conceptual stuff)
A
Most of the matter we encounter consists of mixtures of different substances. Each substance in a mixture retains it scheme identity and properties. The composition of a mixture can vary (unlike substances). The substances making up a mixture are called components of the mixture. Some mixtures do not have the same composition, properties, and appearance throughout. Ex: rocks, hard wood, coffee with sugar
14
Q
Heterogeneous mixture
A
Vary in texture and appearance in any typical sample
15
Q
Homogeneous mixture/solutions
A
Mixtures that are uniform throughout. Solutions can be solids, liquids, or gases. Ex: air
16
Q
Physical properties
A
Can be observed without changing the identity and composition of the substance. Ex: color, odor, density, melting point, boiling point, hardness
17
Q
Chemical properties
A
Describe the way a substance may change/react to form other substances. Ex: flammability
18
Q
Intensive properties
A
Do not depend on the amount of sample being examined and can be used to identify substances. Ex: temperature, melting point
19
Q
Extensive properties
A
Depend on the amount of sample and relate to the AMOUNT of substance present. Ex: mass and volume
20
Q
Physical change
A
A substance changes its physical appearance but not its composition (it is the same substance before and after the change.) All changes of state (changes in temperature and pressure)
21
Q
Chemical change/chemical reaction
A
A substance is transformed into a chemically different substance. Ex: hydrogen burning in air because it combines with oxygen to form water, digesting food, mixing acid and base
22
Q
Distillation
A
Separates the components of a a homogeneous mixture. The process depends on the different abilities of substances to form gases. Ex: boiling a solution of salt and water: water will evaporate and the salt is left behind
23
Q
Chromatography
A
Technique used to separate mixtures based on the differing abilities of substances to adhere to the surfaces to the surfaces of solids
24
Q
Energy
A
The capacity to do work or transfer heat. Does not have a mass but its effects can be observed and measure
25
Work
Energy transferred when a force exerted on an object causes a displacement of that object. w = F*d
F= force exerted on the object
d= distance
26
Heat
The energy used to cause the temperature of an object to increase
27
Force
Any push or pull exerted on the object. Ex: gravity and the attraction between opposite poles of a bar magnet
28
Kinetic energy
Energy of motion. Increases as an objects velocity or speed increases. The transfer of heat is the transfer of kinetic energy at the molecular level
29
Potential energy
Stored energy that arises from the attractions and repulsions an object experiences in relation to tore objects
30
Chemical energy
Released when bonds between atoms are formed, and consumed when bonds between atoms are broken
31
Quantitative
Associated with numbers
32
SI units
Particular choice of metric units for use in scientific measurements. Has 7 base unites from which all other units are derived
33
Length
Meter m
34
Mass
Kilogram kg
35
Temperature
Kelvin K
36
Time
Secondly s or sec
37
Amount of substance
Mole mol
38
Electric current
Ampere A or amp
39
Luminous intensity
Candela cd
40
Law of Conservation Mass
In a chemical reaction, there is no change in the total mass of the materials reacting as compared with the materials that are formed
41
Scientific method
Collect information, formulate a hypothesis, test the hypothesis, formulate a theory, repeatedly test theory
42
Peta
P 10^15
43
Tera
T 10^12
44
Giga
G 10^9
45
Mega
M 10^6
46
Kilo
k 10^3
47
Deci
d 10^-1
48
Centi
c 10^-2
49
Milli
m 10^-3
50
Micro
u 10^-6
51
Nano
n 10^-9
52
Pico
p 10^-12
53
Femto
f 10^-15
54
Atto
a 10^-18
55
Zepto
z 10^-21
56
Mass
A measure of the amount of material in an object. SI base unit of mass is the kg, about 2.2 lbs
57
Temperature
A measure of the hotness or coldness of an object. A physical property that determines the direction of heat flow
58
Heat flow
Always follows from a substance at a higher temperature to one of lower temperature
59
Absolute zero
Zero on the Kelvin scale. The temp at which all thermal motion ceases. -273.15 C
60
Derived unit
Obtained by multiplying or dividing of one or more of the baseunits
61
Volume
Derived SI unit is m^3
62
Density
The amount of mass in a unit volume of a substance
| density=mass/volume
63
Calorie
1 cal = 4.184J
| 1 Cal = 1000 cal = 1 kcal
64
Inexact numbers
Those whose values have some uncertainty. Numbers obtained by measurement
65
Precision
A measure of how closely individual measurements agree with one another
66
Accuracy
Refers to how closely individual measurements agree with the correct or true value
67
Standard deviation
Reflects how much the individual measurements differ from the average
68
Significant figure rules
Zeros between nonzero digits are always significant.
Zeros at the beginning are never significant.
Trailing zeros are significant if the number contains a decimal point
69
Dimensional analysis
Units are multiplied together or divided into each other along with the numerical values. Equivalent units cancel each other
Given unit * desired unit/given unit = desired unit
70
Conversion factor
Fraction whose numerator and denominator are the same quantity expressed in different units
71
Law of constant composition
In a given compound, the relative numbers and kinds of atoms are constant
72
Law of conservation mass
The total mass of materials present after a chemical reaction is the same as the total mass present before the reaction. Atoms are neither created no destroyed during a chemical reaction, the changes that occur during any reaction merely rearrange the atoms
73
Law of multiple proportions
If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers. Ex: ratio of masses of oxygen per gram of hydrogen in water
is 2:1
74
Radioactivity
The spontaneous emission of radiation
75
Cathode rays
Radiation produced between the electrodes when a high voltage was applied to the electrodes in the tube. Originated at the negative electrode and traveled to the positive electrode. Although rays cannot be seen, their presence was detected because they fluoresce (give off light)
76
Thomson's cathode ray experiment
The electron is a negatively charged particle so the electric field deflected the rays in one direction, and the magnetic field deflected them in the opposite direction. Thomson adjusted the strength of the fields so that the effects balanced each other, allowing the electrons to travel in a straight path to the fluorescent screen
77
Rutherford's experiment with 3 types of radiation
3 types of radiation: alpha, beta, and gamma (y). Showed that the paths of alpha and beta radiation are bent by an electric field, in opposite directions. Beta rays bend upwards towards the positively charge place. Gamma rays are unaffected by the plates and go in a straight line. Alpha rays bend downward towards the negative charge plate.
78
What makes an atom of one element different from an atom of another element?
The atoms of each element have a characteristic number of protons
79
Atomic number
The number of protons in an atomy of any particular number. Hydrogen has an atomic number of 1. Also the number of electrons
80
Mass number/atomic weight
Number of protons plus neutrons in the atom. Carbon has an atomic weight of 12.0107amu
81
Isotopes
Atoms with identical atomic numbers but different mass numbers (different number of neutrons)
82
Atomic weight
The average atomic mass of an element. It is found by summing over the masses of its isotopes multiplied by their relative abundances
83
Periodic table
The arrangement of elements in order of increasing atomic number, with elements having similar properties place in vertical columns
84
Groups
Vertical columns. Elements in a group often exhibit similarities in physical and chemical properties because they have the same arrangement of electrons at the periphery of their atoms
85
Metallic elements/Metals
All elements on the left and middle except for hydrogen
86
Nonmetallic elements/metals
Separated from the metals by a stepped line that runs through B, Is, As, Te, At.
87
Metalloids
Elements that lie along the stepped line that has properties that far between those of metals and nonmetals
88
Molecular form
Two or more of the same type of atom bounded together
89
Chemical formula
Ex: O2
90
Molecular compounds
Compounds composed of molecules contain more than one type of atom. Ex: a molecule of the compound CH4. Most molecular compounds are nonmetals
91
Molecular formula
Chemical formulas that indicate the actual numbers of atoms in a molecule Ex: C2H2
92
Empirical formula
Chemical formulas the give only the relative number of atoms of each type in molecule. The subscripts in an empirical formula are always the smallest possible whole-number ratios
Ex: Molecular formula of ethylene is C2H4 and the empirical formula s CH2
93
Structural formula
Depicts atoms represented by their chemical symbols and lines to represent bonds
94
Perspective drawings
Use wedges and dashed lines to depict bonds that are not in the plane of the paper. Give a crude sense of the 3D shape of a molecule
95
Ball and stick models
Show atoms as spheres and bonds as sticks. Accurately represents the angles at which the atoms are attached to one another in a molecule
96
Space filling models
Depict what a molecule would look like if the atoms were scaled up in size. Show the relative size of the atoms. Useful for picturing how two molecules might fit together or pack in the solid state. Doesn't show angles between atoms as well as ball and stick models
97
Ions
Metal atoms tend to lose electrons to formations and nonmetal atoms tend to gain electrons to form anions
98
Ionic compounds
Tend to be composed of a metal cation and nonmetal anion. Ex: NaCl
99
Polyatomic ions
Consists of atoms joined as a molecule but carrying a net positive or negative charge. Ex: NH41+ and SO42-
100
Chemical equations
Represent chemical reactions
101
Combination reactions
Two or more substances react to form one product. A combination reaction between a metaled a nonmetal produces an ionic solid
A+B-->C
102
Decomposition reaction
A single reactant breaks apart to form two or more substances
Many metal carbonates decompose to metal oxides and carbon dioxide when heated
Ex: BaCO3 --> BaO(s) + CO2(g)
C-->A+B
103
Combustion Reaction
Rapid reactions that produce a flame. The combustion of.a hydrocarbon produces CO2 and H2O
H2O gas or liquid depends the reaction conditions
104
Formula Weight
The sum of the atomic weights of the atoms in the chemical formula of the substances in atomic mass units.
If it is an ionic substance, sum the atomic weights of the atoms in the empirical formula
105
Molecular Weight
The formula weight if the chemical formula is that of a molecule
106
Percentage composition/Elemental composition
The percentage by mass contributed by each element in the substance. If the chemical formula is know use the equation on pg 92
107
Mole/Avogadro's number (N_A)
SI Unit for the amount of a substance. The amount of matter that contains the number of atoms in exactly 12 g of isotopically pure 12C which is 6.0221415 x 10^23
108
Molar mass
The atomic weight of an element in atomic mass units is numerically equal to the mass in grams of 1 mol of that element
Ex: Cl has atomic weight of 35.5 amu. 1 mol of Cl has a mass of 35.5 g
The molar mass in grams per mole of any substance is numerically equal to its formula weight in amu
109
Limiting reactant
The reactant that is completely consumed in a reaction. It determines (limits) the amount of the product formed. Once all of the limiting reactant is consumed, the reaction stops
110
Excess reactants
Reactant that is not used up
111
Theoretical yield
The quantity of product calculated to form when all of a limiting reactant is consumed
112
Actual yield
The amount of product actually obtained. Is almost always less than the theoretical yield and can never be greater than it
113
Percent yield
Relates actual and theoretical yields
114
Solvent
Substance present in the greatest quantity
115
Solute
Substance that is dissolved in the solvent
116
Electrolyte
A substance whose aqueous solutions contain ions; any electrolytic solution that conducts an electric current
117
Nonelectrolyte
A substance that does not form ions in solution, usually a molecular compound. A substance that does not ionize in water and consequently gives a nonconducting solution
118
Dissociation
General process in which molecules (or ionic compounds such as salts, or complexes) separate or split into smaller particles such as atoms or ions
Ex: NaCl dissolves in water, each ion separates from the solid structure and disperses throughout the solution: it dissociates into its component ions as it dissolves
119
Solvation
For an ionic compound to dissolve into a liquid, the ions must dissociate and become surrounded by solvent molecules (when water is the solvent it is called hydration)
- In chemical equations, solvated ions are denoted by (aq)
- Helps stabilize the ions in the solution & prevents the cations and anions from recombining
- Ions become dispersed uniformly throughout the solution (ions and their shells of surrounding water molecules are free to move about)
120
Ionize
When a molecule dissociates into its ions in solution
121
Strong electrolytes
Solutes that exist in solution completely or nearly completely as ions.
- Essentially all water-soluble ionic compounds and a few molecular compounds (such as HCl)
122
Weak electrolytes
Solutes that exist in solution mostly in the form of neutral molecules with only small fraction in the form of ions
***Solubility does not indicate whether an electrolyte is strong or weak***
123
Chemical equilibrium
When the relative numbers of each type of ion or molecule in the reaction are constant over time. Represented by half-pointing arrows that go both ways
124
Single reaction arrow
Used for reactions that largely go forward, such as the ionization of strong electrolytes
125
Precipitate
Insoluble product
126
Precipitation reactions
Reactions that results in the formation of a precipitate. Occurs when pairs of oppositely charged ions attract each other so strongly that they form an insoluble ionic solid. Cations and anions come together to form an insoluble ionic compound
Ex: Pb(NO3)2(aq)+2KI(aq)-->PbI2(s)+2KNO3(aq)
127
Solubility
Amount of a substance that dissolves in a given quantity of solvent at a given temperature to form a saturated solution
- Any substance with a solubility less than 0.01mol/L is INSOLUBLE, the attraction between the oppositely charged ions in the solid is too great for the water molecules to separate for the ions to any significant extent and the substance remains largely undissolved
128
Solubility Rules
1) Nearly all nitrates and acetates are soluble
2) All chlorides are soluble except AgCl, Hg2Cl2, PbCl2 (PbCl2 is soluble in hot water)
3) All sulfates are soluble except BaSO4, SrSO4, PbSO4. CaSO4 and Ag2SO4 are only slightly soluble
4) Most of the alkali metal salts and ammonium salts are soluble
5) All the common acids are soluble
6) All oxides and hydroxides are insoluble except those of alkali metals and certain alkaline earth metals
7) All sulfides are insoluble except those of the alkali metals, alkaline earth metals, and ammonium sulfide
8) All phosphates and carbonates are insoluble except those of alkali metals and ammonium salts
129
How to predict whether a precipitate forms?
Note all the ions present in the reactants. Consider the possible cation-anion combinations. Use solubility rules to determine if any of the combinations are insoluble
130
Exchange (Metathesis) reactions
Reaction between compounds that when written as a molecular equation, appears to involve the exchange of ions between the two reactions
Precipitation and acid-base neutralization reactions conform to this pattern
AX+BY-->AY+BX
131
Molecular equation
Shows the complete chemical formulas of reactants and products without indicating ionic character
132
Complete ionic equation
Chemical equation in which dissolved strong electrolytes are written as separate ions
133
Spectator ions
Ions that go through a reaction unchanged and appear on both sides of the complete ionic equation. They can be cancelled on either side of the reaction arrow since they are not reaction with it (once cancelled, the net ionic equation is left)
134
Net ionic equation
Chemical equation for a solution in which soluble strong electrolytes are written as ions and spectator ions are omitted. Only includes the ions and molecules directly involved in the reaction
- Because charge is conserved in reactions, the sum of the ionic charges must be the same on both sides of a balanced net ionic equation
- If every ion in a complete ionic equation is a spectator ion, no reaction occurs
- Net ionic equations illustrate the similarities between various reactants involving electrolytes
Ex: Pb2+(aq)+2I-(aq)-->PbI2(s)
135
Acids
Substances that ionize in aqueous solutions to form hydrogen ions. Because a hydrogen atom consists of a proton and an electron, H+ is simply a proton. Thus acids are called proton donors
136
Monoprotic acids
Yield one H+ per molecule of acid.
| Ex: HCl(aq) and HNO3(aq)
137
Diprotic acids
Yield 2 H+ atoms per molecule of acid. The ionization of diprotic acids occur in 2 steps
Ex: H2SO4(aq)-->H+(aq)+HSO4-(aq)
HSO4-(aq)-->
138
Bases
Substances that accept (react with) hydrogen ions and produce hydroxide ions when they dissolve in water
- Ionic hydroxide compounds such as NaOh,KOH, and Ca(OH)2 are among the most common bases. When dissolved in water, they dissociate into ions, introducing OH- into the solution
- Compounds that do not contain OH- can also be bases: Ammonia (NH3)
139
Strong acids and bases
Acids and bases that are strong electrolytes (completely ionized in solution)
140
Weak acids and bases
Weak electrolytes (partially ionized). The most common weak base is ammonia (NH3)
141
Molecular substance
Any molecular substance that is not an acid or ammonia is probably a nonelectrolyte
142
Neutralization reactions
When a solution of an acid and a solution of a base are mixed. Products of the reaction have none of the characteristic properties of either the acidic solution or basic solution. Protons are transferred from one reactant to another
- In general, a neutralization reaction between an acid and a metal hydroxide produces water and a salt
- Because ions exchange partners, neutralization reactions between acids and metal hydroxides are metathesis reactions
Ex: HCl is mixed with a solution of sodium hydroxide to produce water and table salt
HCl(aq)+NaOH(aq)-->H2O(l)+NaCl(aq)
143
What other bases besides hydroxide ions react with hydrogen ions to form molecular compounds?
Sulfides and carbonate ions. Both of these anions react with acids to form gases that have low solubilities in water
Ex: H2S forms when an acid reacts with a metal sulfide
2HCl(aq)+Na2S(aq)-->H2S(g)+2NaCl(aq)
Carbonates and bicarbonates react with acids to form carbon dioxide gas. Reaction of carbonate or bicarbonate with an acid first gives carbonic acid (H2CO3) which is unstable. If present in solution in sufficient concentrations, It decomposes to water and CO2 gas
144
Oxidation-Reduction/Redox reactions
A reaction in which certain atoms undergo changes in oxidation states. The substance increasing in oxidation state is oxidized; the substance of decreasing in oxidations state is reduced. Electrons are transferred from one reactant to another
145
Oxidized
When an atom, ion, or molecules becomes more positively charged (when it loses electrons)
146
Reduced
When an atom, ion, or molecule becomes more negatively charged (gains electrons)
- When one reactant is oxidized, another reactant must gain them (reduced)
- Oxidation of one substance must be accompanied by the reduction of some other substance
147
Oxidation numbers/states
A way to keep track of electrons gained by the substance being reduced and electrons lost by the substance being oxidized.
- Each atom in a neutral substance of ion is assigned an oxidation number
- For monatomic ions, the oxidation number is the same as the charge
- For neutral molecules and polyatomic ions, the oxidation number of a given atom is a hypothetical charge. This charge is assigned by artificially dividing up the electrons among the atoms in the molecule or ion
148
Displacement reaction/Single Replacement reaction
The reaction between a metal and either an acid or a metal salt. Called displacement reaction because the ion in solution is replaced through oxidation of an element. Many metals undergo displacement reactions with acids, producing salts and hydrogen gas
Conforms to the general pattern:
A+BX-->AX+B
149
Activity series
Metals at the top (alkali metals and alkaline earth metals) are most easily oxidized. They react readily to form compounds and are called active metals. Metals at the bottom of the activity series are very stable and form compounds less readily. These are called noble metals.
***Any metal on the list can be oxidized by the ions of elements below it***
150
Concentration
The quantity of solute present in a given quantity of solvent of solution
- Used to designate the amount of solute dissolved in a given quantity of solvent or quantity of solution
- The greater the amount of solute dissolved in a certain amount of solvent, the more concentrated the resulting solution
151
Molarity
Expresses the concentration of a solution as the number of moles of solute in a liter of solution
M=moles of solute/volume of solution in L
152
Expressing concentrations of an electrolyte
When an ionic compound dissolves, the relative concentrations of the ions in the solution depend on the chemical formula of the compound
Ex: 1.0M solution of NaCl is 1.0M in Na+ and 1.0M in Cl-
1.0M solution of Na2SO4 is 2.0M in Na+ and 1.0M in SO42-
153
Dilution
The process of preparing a less concentrated solution from a more concentrated one by adding solvent
- When solvent is added to the solution, the number of moles of solute remains unchanged:
Moles of solute before dilution=moles of solute after dilution Mconc*Vconc=Mdil*Vdil
154
Titration
Involves combining a solution where the solute concentration is not known with a reagent solution of known concentration (standard solution). Just enough standard solution is added to completely react with the solute in the solution of the unknown concentration. Titrations are carried out to determine the concentration of particular solute in a solution
155
Equivalence point
The point at which stoichiometrically equivalent quantities are brought together
156
How can titrations be conducted?
Neutralization, precipitation, oxidation-reduction
157
Thermodynamics
The study of energy and its transformations
158
Thermochemistry
The relationships between chemical reactions and energy changes that involve heat
159
Energy that originates from chemical reactions
Energy that originates from chemical reactions is associated mainly with changes in potential energy. This energy results from electrostatic interactions between charged particles. Thus, if we are to understand the energy associated with chemical reactions, we must first understand electrostatic potential energy, which arises from the interactions between charged particles
160
Electrostatic potential energy (Eel)
Associated with two charged particles & is proportional to their electrical charges, Q1 and Q2, and is inversely proportional to the distance separating them
Eel=kQ1Q2/d
161
k
Proportionality constant whose value is 8.99 x 10^9 J-m/C^2
162
Joule
Used to measure energy. 1J = 1kg-m^2/s^2
163
Eel values
At finite separation distances, Eel is positive for objects with like charges and negative for objects that are oppositely charged. As the particles move farther apart, their electrostatic potential energy approaches 0
164
Ionic bonding
Based on the electrostatic attraction between cations and anions. To separate ions, we must break the ionic bonds between the ions, which increases the potential energy. The energy to do this must come from some other source. The reverse process in which ions separated by a large distance are allowed to come together to form ionic bonds lowers the potential energy and therefore releases energy
165
Energy and bond formation
Energy is released when chemical bonds are formed.
| Energy is consumed when chemical bonds are broken
166
First law of thermodynamics
Energy can be converted from one form to another, but it is neither created nor destroyed
- To apply this law quantitatively, we need to divide the universe into a finite system of interest to us, and define the energy of that system more precisely
167
System
The portion we single out for a study
168
Surroundings
Everything else
Ex: When we study the energy change that accompanies a chemical reaction in a lab, the reactants and products constitute the system. The container and everything beyond it are considered the surroundings
169
Open system
One in which matter and every can be exchanged with the surroundings
Ex: An uncovered pot of boiling water on a stove: heat comes into the system from the stove and water in released into the surroundings as steam
170
Closed system
Systems that can exchange energy but not matter with their surroundings. These are systems that can be most readily studied under thermochemistry
Ex: A mixture of H2 gas and O2 gas in a cylinder fitted with a piston. The system is the H2 gas and O2 gas. The cylinder, piston, and everything beyond is liberated. Although the chemical form of the H2 and O2 gas atoms in the system is changed in this reaction, the system has not lost or gained mass (it has not exchanged any matter with its surroundings) However, it can exchange energy with its surroundings in the form of work and heat
171
Isolated system
One in which neither energy nor matter can be exchanged with the surroundings
Ex: An insulated thermos containing hot coffee approximates as an isolated system
172
Internal energy (E)
Internal energy of a system is the sum of all the kinetic and potential energies of the components of the system
- The numerical value of a system's internal energy is generally unknown. In thermodynamics, we are mainly concerned with the change in E that accompanies a change in the system
173
Initial internal energy, E final and Delta E
A system with an initial internal energy is represented by Einitial. When the system undergoes a change (which might involve work being done or heat being transferred), the final internal energy of the system after the change is represented by Final. The CHANGE in internal energy is denoted delta E as the difference between the final and initial internal energy:
Delta E = Efinal-Einitial
The initial state of the system refers to the reactants
The final states refer to the products
174
3 parts of the thermodynamic quantity Delta E
A number, a unit, a sign
175
What gives the magnitude of the change for delta E?
The number and unit give the magnitude of the change
176
What gives the direction of delta E?
The sign
177
Delta E values
A positive value of delta E results when Efinal>Einitial. This indicates that the system has gained energy from its surroundings. A negative value of Delta E results when Efinal
178
How can a system exchange energy with its surroundings?
As heat or as work. The internal energy of a system changes in magnitude as heat is added to or removed from the system or as work is done on or by the system
179
When a system undergoes any chemical or physical change
The accompanying change in internal energy (Delta E) is the sum of the heat added to or liberated from the system, q, and the work done on or by the system, w
Delta E=q+w
180
q
The system
+ means the system gains heat
- means the system loses heat
***Not a state function
181
w
Work done on or by the system
+ means work done on system
- means work done by system
***Not a state function
182
Delta E
+ means net gain of energy by system
| - means net loss of energy by system
183
Endothermic
When a process occurs in which the system absorbs heat. Heat flows into the system from its surroundings.
Ex: melting of ice, heat flows into the system from its surroundings
184
Exothermic
A process in which the system loses heat.
| Ex: Combustion of gasoline, heat exits/flows out of the system into the surroundings
185
What conditions influence internal energy?
Temperature and pressure. The internal energy of a system is also proportional to the total quantity of matter in the system because energy is an extensive property
186
State function
A property of a system that is determined by specifying the system's condition or state (in terms of temperature, pressure, etc). The value of a state auction depends only on the present of the system, not the path the system took to reach that state
187
Is E a state function?
Yes. Internal energy is a state function and therefore delta E depends only on the initial and final states of the system, NOT on how the change occurs
188
What does the value of a state function depend on?
The value of a state function depends only on the present state of the system, not on the path the system took to reach that state
189
Enthalpy
Represented by the symbol H and is defined as the internal energy plus the product of the pressure, P, and volume, V of the system
Under conditions of constant pressure, enthalpy provides a function that is a state function and relates mainly to heat flow
H=E+PV
190
Pressure-volume work
Work involved in the expansion or compression of gases. When pressure is a constant in a pressure, the sign and magnitude of the P-V work are given by:
w=-PDeltaV
P is pressure. Always a positive number or 0
DeltaV=Vfinal and Initial is the change in volume of the system
When a gas expands, the system does work on the surroundings and is indicated by a negative value of w
If gas is compressed, deltaV is negative (volume decreases) and w is positive, as work is done on the system by the surroundings
191
1 L-atm
101.3 J
192
Enthalpy change
When a change occurs over constant pressure, the change in enthalpy, deltaH is given by the relationship:
deltaH= delta(E+PV)
=deltaE +PdeltaV (constant pressure)
The change in enthalpy equals the change in internal energy plus the product of the constant pressure ad the change in volume
193
deltaH signs
When deltaH is negative, the system has gained heat from the surroundings, meaning the process is endothermic
When deltaH is positive, the system has released heat to the surrounding, which means the process is exothermic
H is also a state function and depends only on the initial and final states of the system
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Enthalpy of reaction/Heat of reaction
The enthalpy change that accompanies a reaction.
deltaHrxn
When we are given a numerical value for deltaHrxn, we must specify the reaction involved. Ex:
When 2 mol of hydrogen gas burn to form 2 mol H2O(g) at constant pressure, the system releases 483.6 kJ of heat.
deltaH= -483.6 kJ negative sign tells us that it is exothermic
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Thermochemical equations
Balanced equations that show the associated enthalpy change
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Enthalpy change from a chemical reaction formula
deltaH=Hproducts-Hreactants
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Guidelines when using thermochemical equations and enthalpy diagrams
1) Enthalpy is an extensive property
2) The enthalpy for a reaction is equal in magnitude, but opposite in sign, to deltaH for the reverse reaction
3) The enthalpy change for a reaction depends on the states of the reactants and products. It is important to specify the states of reactants and products in thermochemical equations
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Calorimetry
The measure of heat flow. A calorimeter is used to measure heat flow
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Heat capacity (C)
The temperature change experienced by an object when it absorbs a certain amount of heat. The heat capacity of an object is the amount of heat required to raise its temperature by 1 K or 1 degree celsius
- The greater the heat capacity, the greater the heat required to produce a given increase in temperature
- For pure substances the heat capacity is usually given for a specified amount of the substance
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Molar heat capacity Cm
The heat capacity of one mole of a substance
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Specific heat capacity/specific heat Cs
The heat capacity of one gram of a substance
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How can specific heat of a substance be determined?
Experimentally by measuring temperature change, deltaT, that a known mass (m) of a substance undergoes when it gains or loses a specific quantity of heat (q)
Specific heat= quantity of heat transferred/grams of substance*temperature change
Cs=q/mdeltaT
Units will be J/g-K or J/g-C
When a sample absorbs heat (positive q), its temperature increases (positive deltaH):
q=CsmdeltaT
We can calculate the quantity of heat a substance gains or loses by using its specific heat together with its measuredness and temperature change
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qsoln
Specific heat*grams of solution*deltaT=-qrxn
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Bomb calorimeter
Used for combustion (when a compound reacts with excess oxygen). The substance is placed in a small cup (the bomb) within an insulated sealed vessel. The heat released when combustion occurs is absorbed by the water and the various components of the calorimeter, causing the temperature of the water to rise. The change in water temperature by the reaction is then measured
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Ccal
The total heat capacity of the calorimeter. This must be known to calculate the heat of combustion from the measured temperature increase. This quantity is determined by combusting a sample that releases a known quantity of heat and measuring the temperature change
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Hess's Law
States that if a reaction is carried out in a series of steps, deltaH for the overall reaction equals the sum of the enthalpy changes for the individual steps
- The overall enthalpy change for the process is independent of the number of steps and independent of the path by which the reaction is carried out
- Law is a consequence that enthalpy is a state function
- deltaH can be calculated for any process as long as we find a route for which deltaH is known for each step
- Provides a useful means of calculating energy changes that are difficult to measure directly
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Enthalpy of formation/heat of formation
deltaHf is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms. The subscript f indicates that the substance has been formed from its constituent elements
- Magnitude of any enthalpy change depends on temperature, pressure, and state of the reactants and products
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Standard state
Used to compare enthalpies of different reactions. It is a defined set of conditions at which most enthalpies are tabulated
- Standard enthalpies of formation are measured under standard conditions (25 degree C and 1.00 atm pressure)
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Standard enthalpy of change
Of a reaction is defined as the enthalpy change when all reactants and products are in their standard states
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Standard enthalpy of formation
the change in enthalpy for the reaction that forms one ole of the compound from its elements with all substances in their standard states:
elements-->compound ( 1 mole in standard state)
deltaHrxn=deltaHfdegree
- If an element exists in more than one form under standard conditions, the most stable form of the element is usually used for the formation reaction
- Reported in kJ/mol of the substance being formed
- Standard enthalpy of formation of the most stable form of any element is 0 because there is no formation reaction needed when the element is already in its standard state
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Standard enthalpy of change of a reaction formula
The sum of the standard enthalpies of formation of the products minus the standard enthalpies of formation of the reactants:
deltaHrxndegree={ndeltaHfdegree(products)-{mdeltaHfdegree(reactants)
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Bonds and chemical energy
The energy changes that accompany chemical reactions are closely related to the changes associated with forming and breaking chemical bonds
- Breaking bonds requires energy
- Forming bonds releases energy
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Bond enthalpy
The enthalpy change deltaH, for the breaking of a particular bond in one mole of a gaseous substance
Ex: Cl2 has a bond enthalpy of 242 kJ/mol. Thebond enthalpy is a positive number because energy must be supplied from the surroundings to break the Cl-Cl bond
***Always a positive quantity when bonds are broken (energy is required to break chemical bonds) ENDOTHERMIC
***When a bond forms, energy is released so negative. EXOTHERMIC
***The greater the bond enthalpy, the stronger the bond
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Enthalpy. If deltaH is > 0
Reaction will be endothermic
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Enthalpy. If dealt is < 0
Reaction will be exothermic
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Bond enthalpy guidelines
1) We supply enough energy to break those bonds in the reactants that are not present in the products. The enthalpy of the system is increased by the sum of the bond enthalpies of the bonds that are broken
2) We form the bonds in the products that were not present in the reactants. This step releases energy and therefore lowers the enthalpy of the system by the sum of the bond enthalpies of the bonds that are formed
deltaHrxn = [sum of bond enthalpies of bonds broken]-[sum of bond enthalpies of bonds formed]
if deltaH is positive, reaction is endothermic
if deltaH is negative, reaction is exothermic
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Electromagnetic radiation/radiant energy
The light we see with our eyes is an example
- Carries through space so it is also known as radiant energy
- Other types besides visible light includes radio waves, infrared radiation (heat), x-rays
- All types move through a vacuum at 2.998 x 10^8 m/s which is known as C, the speed of light
- All have wavelike characteristics similar to those of waves that move through water
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Periodic (wave nature of light)
Patterns of peaks and troughs repeat itself at regular intervals
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Wavelength
Lambda, distance between two adjacent peaks or troughs
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Frequency
v, The number of wavelengths (cycles) that pass a given point each second
- Expressed in cycles per second (a unit called hertz, Hz)
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Inverse relationship between the frequency and wavelength of electromagnetic radiation
lambda is wavelength
v is frequency
c is speed of light
lambda*v=c
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Hertz
Hz, frequency expressed in cycles per second. Usually given as per second denoted by /s or s^-1
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Observations that cannot be resolved by the wave model of light (3 of them)
1) The emission of light from hot objects (known as blackbody radiation because the objects studied appear black before heating)
2) The emission of electrons from metal surfaces on which light shines (the photoelectric effect)
3) The emission of light from electronically excited gas atoms (emission spectra)
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What does the wavelength distribution of the radiation depend on?
Temperature.
| Ex: a red-hot object is cooler than a yellow or white one
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Planck's theory and constant
E=hv
h is a constant 6.626 x 10^-34 J-s
- According to Planck's theory, matter can emit and absorb energy only in whole number multiples (1hv, 2hv, 3hv)
- If the quantity of energy emitted by an atom is 3hv, the three quanta of energy have been emitted
- Energy can be released only in specific amounts, so the allowed energies are quantized (their values are restricted to certain quantities)
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Quantum
Fixed amount
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Photoelectric effect
Albert Einstein used Planck's theory to explain the photoelectric effect.
- Light shining on a clean metal surface causes electrons to be emitted from the surface.
- A minimum frequency of light, different for different metal, is required for the emission of electrons
Ex: light with a frequency of 4.60 x 10^14/s or greater causes cesium metal to emit electrons, but if the light has frequency less than that, no electrons are emitted
- Einstein assumed that the radiant energy striking the metal surface behaves like a stream of tiny energy packets. Each packet is a photon
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Photon
Particle of energy.
- Einstein deduced that each photon must have an energy equal to the Planck constant times the frequency of light:
Energy of photon = E = hv
- Einstein's theory of light as a stream of photons. Light possesses both wave-like and particle-like characteristics
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h
Planck's constant
| 6.626 x 10^34 J-s
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Work function
Certain amount of energy
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Photons and the photoelectric effect
Under the right conditions, photons striking a metal surface can transfer their energy to electrons in the metal. A certain amount of energy, the work function, is required for the electrons to overcome the attractive forces holding them in the metal.
* **Increasing the intensity of the light source doesn't lead to emission of electrons from the metal; only changing the frequency of the incoming light has that effect. The intensity (brightness) of the light is related to the number of photons striking the surface per unit time but not to the energy of the photon.
- When the frequency is such that photons have energy greater than the work function of the particular metal, electrons are emitted
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Monochromatic
Radiation composed of a single wavelength
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Polychromatic radiation
Radiation containing many different wavelengths
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Spectrum
Produced when radiation from a polychromatic source is separated into its component wavelengths. The resulting spectrum consists of a continuous range of colors
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Continuous spectrum
Rainbow of colors, containing light of all wavelengths
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Line spectrum
A spectrum containing radiation of only specific wavelengths
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Rydberg equation
Allows us to calculate the wavelengths of all the spectral lines of hydrogen
Rydberg constant is 1.096776 x 10^7/m
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Principal quantum number
n. Can have whole number values of 1,2,3,etc
- Describes an orbit
- Each allowed orbit corresponds to a different value of n. The radius of the orbit gets larger as n increases
The first orbit, closest to the nucleus n=1, the next n=2, and so on
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Ground state
n=1
| The lowest energy state of the atom
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Excited state
n-2 or higher. This is when the electron is in a higher-energy state
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Bohr's 3 postulates
1) Only orbits of certain radii, corresponding to certain specific energies, are permitted for the electron in a hydrogen atom
2) An electron in a permitted orbit is in an allowed energy state. An electron in an allowed energy state does not radiate energy and therefore, does not spiral into the nucleus
3) Energy if admitted or absorbed by the electron as the electron changes from one allowed energy state to another. this energy is emitted or absorbed a a photon that has energy e=hv
- for part 3) electrons jumps from one allowed orbit to another by either absorbing or emitting photons whose radiant energy corresponds exactly to the energy difference between the two orbits. Energy is absorbed to move to a higher n value (high energy state) and energy is emitted to jump to a lower energy state (lower n value)
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Ephoton
The energy of the photon (Ephoton) must equal the difference in energy between the two states, deltaE
deltaE positive = energy is absorbed
deltaE negative = energy is emitted
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Bohr model limitiation
Can't explain the spectra of other atoms besides hydrogen
- Bohr avoided the problem of why the negatively charged electron would not just fall into the positively charged nucleus, by assuming it would not happen
3) Electrons orbiting the nucleus at a fixed distance is not a realistic model
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2 important ideas we got from Bohr's model
1) Electrons only exist in. certain discrete energy levels, which are described by quantum numbers
2) Energy is involved in the transition of an electron from one level to another
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De Broglie
Suggested that an electron moving about the nucleus of an atom behaves like a wave and therefore has a wavelength. Proposed that wavelength of the electron depends on its mass and velocity
wavelength = h/mv
mv is the momentum (mass*velocity)
h is Planck's constant
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Matter waves
De Broglie used this term to describe the wave characteristics of material particles
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Uncertainty principle
Werner Heisenberg proposed that the uncertainty principle: it is impossible for us to know simultaneously both the exact momentum of an electron and its exact location in space
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Wave functions
Functions that describes the electron in the atom. these functions are represented by Greek lowercase letter psi
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psi/probability density/electron density
Psi squared, at a given point in space represents the probability that the electron will be found at that location`
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Orbitals
Set of wave functions. Each orbital has a characteristic shape and energy
***DO NOT CONFUSE WITH ORBIT. An orbit visualizes the electron moving in a physical orbit. An orbital is a quantum-mechanical model which describes electrons in terms of probabilities, visualized as electron clouds
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Quantum-mechanical model
Uses three quantum numbers: n, l, ml
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Angular momentum quantum number
l, the second quantum number. Can have integral values of 0 to n-1 for each value of n. This defines the shape of the orbital. The value of l for a particular orbital is generally designated by the letters s, p ,d,f corresponding to l values of 0, 1,2,3
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Magnetic quantum number
ml. Can have integral values between -l and l, including 0. This quantum number describes the orientation of the orbital in space
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Electron shell
The collection of orbitals with the same value of n
| Ex: all orbitals that have n=3 are said to be in the third shell
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Subshell
The set of orbitals that have the same n and l values. Each sub shell is designated by a number (value of n) and a letter (s,p,d,f, corresponding to the value of l)
Ex: orbitals that have n=3 and l=2 are called 3d orbitals and are in the 3d subshell
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s Orbital
All s orbitals are spherically symmetric. The l quantum number for the s orbital is 0 and therefore, the ml quantum number is also 0. Thus, for each value of n, there is only one s orbital.
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Radial probability density
The probability that the electron is at a specific distance from the nucleus
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p Orbitals
l quantum number for the p orbital is 1. Each p sub shell has 3 orbitals, corresponding to the 3 allowed values of ml: -1,0,1. Electron density is not distributed spherically as in an s orbital, it is dumbbell shaped and has 2 lobes
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d Orbital
When n is 3 or greater, we encounter d orbitals (l=2). There are five 3d orbitals. Each shell has 5 possible values for the ml quantum number: -2,-1,0,1,2. D orbitals have a 4 leaf clover shape with 4 lobes.
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f Orbital
When n if 4 or greater, there are 7 equivalent f orbitals (l=3)
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Degenerate
Orbitals with the same energy
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Electron spin
An intrinsic property that causes each electron to behave as if it were a tiny sphere spinning on its own axis
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Spin magnetic quantum number
ms. Can be +1/2 or -1/2
An upwards half arrow is a +1/2. Down is -1/2
A spinning charge produces a magnetic field. The two opposite directions of spin therefore produce oppositely charge magnetic fields. these two opposite magnetic fields lead to the splitting of spectral lines into close spaced pairs.
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Pauli exclusion principle
No two electrons in an atom can have the same set of four quantum numbers n, l, ml, and ms. For a given orbital, the values of n, l and ml are fixed. Thus, if we want to put more than one electron in an orbital AND satisfy the Pauli exclusion principle, the only choice is to assign different ms values to the electrons. Because there are only 2 values, we conclude that an orbital can hold a maximum of 2 electrons and must have opposite spins
- Orbitals are filled in order of increasing energy, with no more than 2 electrons per orbital
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Electron configuration
The way electrons are distributed among the various orbitals of an atom
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Ground state
The most stable electron configuration. The one in which the electrons are in the lowest possible energy states
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Orbital diagram
Diagram in which each orbital is denoted by a box and each electron by a half arrow. Pictorials of the electron spin corresponds to the directions of the magnetic fields.
- Electrons having opposite spins are paired
- Unpaired electrons are not accompanied by a partner of opposite spin
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Hund's Rule
States that when filling degenerate orbitals the lowest energy is attained when the number of electrons having the same spin is maximized
* **Electrons occupy orbitals singly to the maximum extent possible and that these single electrons in a given sub shell all have the same spin magnetic quantum number. Electrons arranged in this way are said to have parallel spins
- Hund's rule is based on the fact that electrons repel one another. By occupying different orbitals, the electrons remain as far as possible from one another, minimizing electron-electron repulsions
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Anomalous electron configurations
Chromium is [Ar]3d54s1
| Copper is [Ar]3d104s1
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Valence bond theory
Bonding electron pairs are concentrated in the regions between atoms, and nonbonding electron pairs lie in directed regions of space.
- The buildup of electron density between two nuclei occurs when a valence atomic orbital of one atom shares space, or overlaps, with the valence atomic orbital of another atom. The overall of orbitals allows two electrons of opposite spin to share the space between the nuclei, forming a covalent bond
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Hybrid orbitals
To explain molecular geometries, we often assume that the atomic orbitals on an atom (usually the central atom) mix to form new orbitals
- Shape of any atomic hybrid orbital is different from the shape of the original atomic orbitals
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Hybridization
The process of mixing atomic orbitals is a mathematical operation
- The total number of atomic orbitals on an atom remains constant, so the total number of hybrid orbitals on an atom equals the number of atomic orbitals that are mixed
- Each hybrid orbital is equivalent to the others but point in opposite directions
- Used to describe the bonding in molecules containing nonbonding pairs of electrons (Ex:H2O)
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sp Hybrid orbitals
According to the valence-bond model, a linear arrangement of electron domains implies sp hybridization
Ex: BeF2 has 2 sp hybrid orbitals, HbCl2
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sp2 Hybrid orbitals
Mixing one 2s and one 2p atomic orbital yield two equivalent sp hybrid orbitals that point in opposite directions
Ex: BF3. Mixing the 2s and two of the 2p atomic orbitals yields 3 equivalent sp2 hybrid orbitals
Ex: SO3 has 3 sp3 hybrid orbitals
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sp3 Hybrid orbitals
An s atomic orbital can mix with all three p atomic orbitals in the same sub shell.
Ex: CH4 the carbon atom forms four equivalent bonds with the four hydrogen atoms. We envision this process as resulting from the mixing of the 2s and all three 2p atomic orbitals of carbon to create four equivalent sp3 hybrid orbitals
Ex: CH4, NH3, H2O, NH4+
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How to describe the hybrid orbitals used by an atom in bonding
1) Draw the lewis structure for the molecule or ion
2) Use the VSEPR model to determine the electron-domain geometry around the central atom
3) Specify the hybrid orbitals needed to accommodate the electron pairs based on their geometric arrangement
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Molecular orbital theory
A more sophisticated model that explains some aspects of bonding. Describes the electrons in molecules by using specific wave functions
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Molecular orbital (MO)
A specific wave function used to describe electrons in molecules
- Similar to atomic orbitals
- Hold up to 2 electrons of the opposite spin
- Has a definite energy
- We can visualize its electron-density distribution by using a contour representation
* ** MO are associated with an entire molecule, not with a single atom
- Whenever two atomic orbitals overlap, two molecular orbitals form. Ex: the overlap of the 1s orbitals of two hydrogen atoms to form H2 produces two MOs
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Constructive combination
Formed by adding the wave functions
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Bonding molecular orbital
The energy of the constructive combination MO is lower in energy than the two atomic orbitals from which it was made
- the electron density is concentrated in the region between the 2 nuclei
- Lower in energy because its more stable
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Destructive combination
mining the two atomic orbitals in a way that causes the electron density to be canceled in the central region where the two overlap
The energy is referred to as the anti bonding molecular orbital, is higher than the energy of the atomic orbitals
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Bond order
In MO theory, the stability of a covalent bond is related to its bond order, defined as half the difference between the number of bonding electrons and the number of anti bonding electrons
Bond order=1/2(number of bonding electrons-number of anti bonding electrons)
- A bond order of 1 represents a single bond, bond order of 2 represents a double bond, so on
