Final Exam Flashcards
(221 cards)
1
Q
what something is made of
A
composition
2
Q
What something is made of and how the components are arranged
A
Structure
3
Q
The fundamental unit of matter
A
Atom
4
Q
Made of one type of atom
A
Element
5
Q
Composed of more than one element, bound in a fixed ratio
A
Compound
6
Q
Contain more than one substance, not bound in a fixed ratio
A
Mixtures
7
Q
Groups of atoms that bind tightly together, and behave as a single unit
A
Molecules
8
Q
Mixed evenly throughout the substance
A
Homogeneous mixture
9
Q
Example of homogeneous mixture
A
Salt mixes evenly with water
10
Q
Not mixed evenly throughout the substance
A
Heterogeneous mixture
11
Q
Example of heterogeneous mixture
A
Sand and water
12
Q
Definite shape and definite volume
A
Solid
13
Q
Definite volume but not definite shape
A
Liquid
14
Q
No definite shape or volume
A
Gas
15
Q
Can be measured without changing the identity of the substance
A
Physical properties
16
Q
Examples of physical properties
A
mass, volume, temperature, color, hardness
17
Q
Don’t change the identity of the substance
A
Physical changes
18
Q
Example of physical change
A
Remolding it, phase change
19
Q
Can not be measured without changing the identity of the substance
A
Chemical properties
20
Q
Example of chemical properties
A
Catching something on fire
21
Q
Change the identity of the substance
A
Chemical change (AKA chemical reactions)
22
Q
Ability to do work
A
Energy
23
Q
Energy that is stored
A
Potential energy
24
Q
The energy of motion
A
Kinetic energy
25
Involves the kinetic energy of the particles in a substance
Heat energy
26
Releases heat energy
Exothermic change
27
Absorbs heat energy
Endothermic change
28
A tentative explanation that has not been tested
Hypothesis
29
An idea supported by experimental evidence
Theory
30
A statement that describes observations that are true in widely varying circumstances
Scientific law
31
-How reliable are the measurements
-Do they reflect the true value
Accuracy
32
-How finely are the measurements made
-How closely are they grouped together
Precision
33
Indicate how precisely we know a measurement
Significant digits
34
Density formula
d=m/v
35
Boiling point, freezing point, and absolute value of water in celsius
Boiling point: 100 C
Freezing point: 0 C
Absolute Value: 273 C
36
common volume units
cm^3= mL
dm^3= L
m^3
37
Convert Celsius to Kelvin
K= C + 273
38
First proposed atoms
Democritus
39
Who proposed law of conservation of Mass
Antoine Lavoisier
40
Law of Conservation of Mass
In chemical reactions, matter is neither created nor destroyed
41
Who proposed atomic theory
John Dalton
42
Who put the periodic table together
Dmitri Mendeleev
43
Who discovered the electron
J. J. Thomson
44
Who showed that atoms are mostly empty space with a small, dense nucleus
Rutherford
45
Who proposed the Bohr model
Bohr
46
The number of protons in an atom
Atomic number
47
The number of protons and neutrons in an atom
Mass number
48
Have the same atomic number, but different mass number
Isotope
49
Where are the metals on the periodic table?
Left two columns except for hydrogen
50
Where are the nonmetals on the periodic table?
The right side and hydrogen
51
Where are the metalloids on the periodic table?
Stairsteps
- B, Si, Ge, As, Sb, Te, Po, At
52
Three fundamental ideas
1. All matter is composed of atoms
2. The atoms of each element have unique characteristics and properties
3. In chemical reactions, atoms are not changed, but combine in whole-number ratios to form compounds
53
Group 1A on the periodic table
Alkali Metals
54
Group 2A on periodic table
Alkaline earth metals
55
Group 7A on periodic table
Halogens
56
Group 8A
Noble Gases
57
the length of one wave
wavelength
58
The number of waves per second
Frequency
59
How is wavelength and frequency related?
They are inversely proportional
60
Observe colors emitted by different metal ions
Flame test
61
What color is produced when an electron falls to a lower orbital?
3->2: Red
4->2: light blue
5->2: indigo (deep blue)
6->2: purple violet
2->1: ultra violet (UV)
62
Describes electrons
Quantum Mechanics
63
How many electrons fit in levels 1 2 3 and 4?
1: 2
2: 8
3: 18
4: 32
64
What are the sublevels and how many orbitals are in each sublevel? and how many electrons are in each orbital?
S: 1 orbital: 2 electrons
P: 3 orbitals: 6 electrons
D: 5 orbitals: 10 electrons
F: 7 orbitals: 14 electrons
65
If empty orbitals of the same energy are available, electrons singly occupy orbitals rather than pairing together
Hund's Rule
66
The highest occupied electron energy level
Valence Level
67
An atom is stabilized by having its highest-occupied (valence) energy level filled
Octet Rule
68
Atoms or groups of atoms that have an overall charge
Ions
69
Positively charged ions
Cations
70
Negatively charged ions
Anions
71
Groups of atoms with a charge
Polyatomic ions
72
An attraction between oppositely charged ions
Ionic bond
73
Composed of charged ions
Ionic compound
74
An array of positive and negative ions
Ionic lattice
75
Show the type and amount of each element present
Chemical formulas
76
The smallest whole-number ratio of atoms
Empirical formula
77
Electrons shared between two atoms
Covalent Bond
78
Gives the number of atoms in the molecule
Molecular formula
79
A homogeneous mixture, in which the main component is water
Aqueous solution
80
Ionic compounds dissolved in water
Electrolyte solutions
81
Ions are pulled apart in an aqueous solution
Dissociation
82
Covalent compounds that produce H+ ions in aqueous solution
Acids
83
Shows ions together as compounds
Molecular equation
84
Shows dissociated ions as separate species
Ionic equation
85
What kind of reaction:
One forms two are more
Decomposition
86
What kind of reaction:
Two form one
Synthesis
87
What kind of reaction:
One element replaces another
Single displacement
88
What kind of reaction:
Two ions replace each other
Double displacement
89
Factors affecting solubility
-Charge on ions
-Size of ions
-How tightly ions pack together
90
2 aqueous solutions produce an insoluble product
Precipitation reaction
91
The solid product formed in the reaction
Precipitate
92
Shows all ions present
Complete ionic equation
93
Only include ions involved in the precipitation
Net ionic equation
94
Shows neutral compounds
Molecular equation
95
Compounds that produce OH- ions in aqueous solution
Bases
96
Reactions in which oxygen gas combines with elements or compounds to produce oxides
Combustion reactions
97
Compounds composed of hydrogen and carbon
Hydrocarbons
98
what do the combustion of hydrocarbons produce
water and carbon dioxide
99
The determination of the proportions in which elemements or compounds react with one another
Mass stoichiometry
100
The mass of a single molecule or formula unit
Formula mass
101
formula for percent composition of one element
mass of 1 element/ mass of compound X 100%
102
how many particles equals one mole?
6.02 X 10^23 particles = 1 mole
103
Using the amount of one material to predict the amount of another, based on the balanced equation
Stoichiometry
104
Completely consumed; limits the amount of product formed
Limiting reagent
105
Not completely consumed; regent will be left over after the reaction is complete
Excess Reagent
106
The amount of a product that can form, based on the balanced equation
Theoretical yield
107
The amount actually obtained
Actual Yield
108
The percent of the theoretical yield that was obtained
Percent Yield
109
How to find percent yield
actual yield/ theoretical yield X 100%
110
Why is actual yield so low?
- Material sticks to container walls
- Unwanted side products
- Product lost during purification
111
The transfer of kinetic energy
Heat
112
The transfer of energy from one form to another
Work
113
The part of the universe being studied
System
114
The rest of the universe
Surroundings
115
Releases heat energy
Exothermic change
116
Absorbs heat energy
Endothermic change
117
Energy cannot be created or destroyed
Law of conservation of energy
118
The total kinetic energy transformed from one substance or object to another
Heat
119
The average kinetic energy of the particles in a substance
Temperature
120
The amount of heat required to raise the temperature of 1 gram of material 1 degree Celsius
Specific heat
121
Specific heat formula
specific heat= heat/ (mass) (change in temperature)
122
The amount of heat required to raise an object by 1 degree celsius
Heat capacity
123
Measure the flow of heat
Calorimetry experiments
124
Measures energy content (heat of reaction) in food and fuels
Bomb calorimetry
125
The energy of a reaction
extensive property
126
The amount of energy that can be produced by the combustion of a material
Fuel Value
127
The amount of heat energy absorbed or released in a chemical reaction at constant pressure
Reaction Enthalpy
128
How to find formal charge
Formal charge= valence electrons in the neutral atom- number of covalent bonds- number of unshared electrons
129
- A set of structures that show how electrons are distrusted
- Used when a single Lewis structure is insufficient
Resonance structure
130
Arrangement of electrons around the central atom
Electronic geometry
131
Shape caused by the arrangement of atoms
Molecular geometry
132
Atoms do not share the electrons evenly
Polar covalent bond
133
How strongly atoms pull bonded electrons
Electronegativity
134
Electronegativity and how it related to the periodic table
- Electronegativity increases going left to right
- Electronegativity increase when going bottom to top
135
Polar covalent scale
Covalent: 0-0.4
Polar covalent: 0.5-2.0
Ionic: 2.1 - +
136
An overall polarity in a molecule
Molecular dipole
137
Particles are close together and held in a fixed place
Solid
138
Definite shape and volume
Solid
139
Particles are closer together but move freely past each other
Liquid
140
Definite volume; adopts the shape of the container
Liquid
141
Particles are far apart and have very little interaction
Gas
142
Adopts shape and volume of container
Gas
143
A transition from one state of matter to another
Phase change
144
Rigid frameworks of atoms, molecules, or ions
Lattices
145
Properties of metallic substances
- Form lattices of tightly packed atoms
- Electrons move easily between atoms
- Shapes of metals are easily altered
146
Lattices of covalent bonds that form giant molecules
Covalent networks
147
Contain long chains of covalently bonded atoms
Polymers
148
Forces within molecules
Covalent bonds
149
Forces between molecules
Intermolecular forces
150
Three types of intermolecular forces
- dipole-dipole interactions
- Hydrogen bonds
- Dispersion forces
151
Attractions between polar covalent molecules
Dipole-dipole interactions
152
A strong intermolecular force between molecules containing H-F, H-O, or H-N bonds
Hydrogen bonding
153
Weak intermolecular forces that result from instantaneous dipoles
Lond Dispersion Forces
154
Strongest to weakest intermolecular forces
1. Hydrogen bonding
2. Dipole-dipole forces
3. London dispersion forces
155
The force that gases exert on their surroundings
Pressure
156
A device used to measure atmospheric pressure
Barometer
157
The difference between the compressed gas pressure and the atmospheric pressure
Gauge pressure
158
The pressure and volume of a gas are inversely related
Boyle's Law
159
Boyle's Law equation
P1V1=P2V2
160
At constant pressure, the volume of a gas is directly proportional to its temperature
Charle's Law
161
Charle's Law formula
V1/T1= V2/T2
162
Combined Gas Law Formula
P1V1/T1= P2V2/T2
163
If temperature and pressure are constant, the volume of a gas is proportional to the number of moles of gas present
Avogadro's Law
164
What does PV=nRT mean
P= Pressure (must be in atm)
V= Volume (must be in L)
R= 0.0821(given)
T= Temperature (must be in Kelvin)
n= number of moles
165
The pressure caused by one gas in a mixture
Partial pressure
166
The spread of particles through random motion
Diffusion
167
The process of a gas escaping from a container
Effusion
168
Lighter particles escape ____
Faster
169
Heavier particles escape _____
Slower
170
The substance that dissolves
Solute
171
The major component of the solution
Solvent
172
The amount of solute present in a solution
Concentration
173
How to find mass%
mass of solute/ mass of solution X 100%
174
How to find volume %
volume of solute/ volume of solution X 100%
175
How to find mass/volume %
mass of solute/ volume of solution X 100%
176
How to find in ppm
mass of solute/ mass of solution X 10^6
177
How to find ppb
mass of solute/ mass of solution X 10^9
178
Molarity formula
M= moles of solute/ liter of solution
179
Preparing dilute solutions
MiVi=MfVf
M=molarity
V=volume
180
The three colligative properties
- Freezing point depression
- Boiling Point elevation
- Osmotic pressure
181
Boiling point elevation
More molecules, the higher the boiling point
182
Freezing point depression
As you increase the solute in concentration the lower the freezing point becomes
183
Water can pass through it, but molecules and ions cannot
Semipermeable membrane
184
Water moves toward the more concentrated solution
Osmotic pressure
185
Water leaves to go to the salt
Hypertonic
186
Water comes to the cell, which has the salt
Hypotonic
187
The same amount of water is leaving as it is coming
Isotonic
188
An ion of one metal reacts with the elemental form to another metal
Metal displacement reactions
189
Acids and bases combine to form what?
water and salt
190
Arrhenius definition of acid
Produce H+ ions in water
191
Arrhenius definition of base
Produce OH- ions in water
192
Bronsted-Lowry definition of acid
Compounds that donate H+ ions
193
Bronsted-Lowry definition of base
Compounds that accept H+ ions
194
Completely ionize in water
Strong acids
195
Partially ionize in water
Weak acids
196
Goes in both forward and backward directions
Equilibrium reaction
197
Can release more than one H+ into aqueous solution
Polyprotic acid
198
How to find the conjugate base
Conjugate base is identical to the acid but without the H+
199
What do metals produce when they react with an acid?
Metal cations and hydrogen gas
200
Nonmetals bonded to oxygen
- Often react with water to form acids
Nonmetal oxides
201
Describes changes that involve the nucleus of the atom
Nuclear chemistry
202
Particles in the nucleus (protons and neutrons)
nucleons
203
Protons only
Atomic number
204
Protons and neutrons
mass number
205
An atom or nucleus containing a particular number of protons and neutrons
Nuclide
206
Have the same atomic numbers but different mass numbers
Isotopes
207
Change in the structure of a nucleus
Nuclear Reactions
208
Show nuclear changes
Nuclear equations
209
The spontaneous release of particles and/or energy (radiation) from a nucleus
Radioactivty
210
Emission of a helium nucleus
Alpha Decay
211
Emission of an electron
Beta decay
-Mass number does not change but atomic number increases by one
212
Release of energy (gamma rays)
Gamma Decay
-Neither the atomic number nor mass number changes
213
A naturally occurring sequence of radioactive decays
Radioactive Decay Series
214
The amount of time required for one-half of a sample of a radioactive substance to decay into something else
Half-Life
215
A powerful force that holds the nucleus together
Nuclear force
216
Difference between the masses of the particles and the nucleus
Mass defect
217
Mass can be converted into energy, and energy can be converted into mass
Mass-energy equivalence
218
Released when the protons and neutrons form a new nucleus
Blinding energy
219
Einstein's equation
E=mc^2
E=Energy
m=mass
c=speed of light
220
A large nucleus shatters into smaller nuclei
Fission
221
Can go under fission
Fissle
