1.1 General Chemistry Flashcards
(168 cards)
1
Q
SOLID (shape & volume)
A
Definite *non-compressible
2
Q
SOLID (Molecular motion)
A
Vibration *2 stones
3
Q
LIQUID (shape)
A
Indefinite *assumes container shape
4
Q
LIQUID (volume)
A
Definite
5
Q
LIQUID (molecular motion)
A
Gliding *ex. water falls
6
Q
GAS (shape & volume)
A
Indefinite *compressible
7
Q
GAS (molecular motion)
A
constant random
8
Q
Plasma/Ionized Gas
A
4th state; most abundant state of matter
• Has p+ and e- (thus, greatly affected by magnetic field)
• Ex. ionized Ne light, Aurora, Stars, Sun
9
Q
IFA Strength most ↑ or strongest:
A
S > L > G > P
10
Q
Enthalpy (heat/ reaction energy):
A
P > G > L > S
11
Q
Melting
A
*aka: Fusion, Liquefaction, Thawing
Solid to Liquid
12
Q
Freezing
A
Liquid to Solid
13
Q
Evaporation
A
Liquid to Gass
14
Q
Condensation
A
Gas to Liquid
15
Q
Sublimation
A
*moth/naphthalene balls
Solid to Gas
16
Q
Deposition
A
*dry ice/cardice
Gas to Solid
17
Q
Recombination
A
*aka: Deionization
Plasma to Gas
18
Q
Ionization
A
Gas to Plasma
19
Q
Pure substance
A
- Element
* Compound
20
Q
Element
A
simplest form of substance
21
Q
Compound
A
2 or more chemical united (separated via
chemical means)
22
Q
Mixture
A
- Homogeneous
* Heterogeneous
23
Q
Mixture
A
2 or more substance wherein individual substance identifies are retained (separated via physical means. Alcohol +
Water via distillation)
24
Q
Homogeneous
A
1 phase; solution *clear colored
25
Heterogeneous
2 phases; suspension, colloid *ex. milk
26
Extrinsic Property “Dependent”
Length, mass/weight, volume, pressure, entropy, enthalpy, electrical resistance
27
Intrinsic Property “Independent”
Density/ SpGr (water = 1g/ml or cc), viscosity (resistance to flow), velocity (m/sec), temperature, color
28
Law of Conservation of Mass/ Matter
- Antoine Lavoiser
| - Mass/ Matter is always constant (neither created nor destroyed)
29
Law of Definite/Constant Proportions
- Joseph Proust (Proust's law)
• Chemical compounds always contain the exact proportion of element in fixed ratio (by mass)
• Ex. H2O →2H + O, C6H12O6 = CH2O
30
Law of Multiple Proportion
```
- John Dalton
• When 2 elements form more than 1 compounds, it can be expressed in a fixed whole number (by mass)
• Ex. CO → 28g/mole, CO2 → 44g/mole
C = 12g/mole
O = 16g/mole
```
31
Law of combining weights
• Proportions by weight when chemical reaction takes place can be expressed in small integral unit
• Ex. MgO → 40g/mole (100%)
Mg = 24g/mole (60%); O = 16g/mole (40%)
32
Democritus
- atomos
| - "indivisible"
33
John Dalton
- Billiard ball
| - Matter is made up of atoms
34
Postulates
* Elements are composed of indivisible, indestructible atoms
* Atoms alike for a given element (isotopes)
* Atoms of different elements differ in size, mass & other properties (isobars)
* Compound are formed form 2 or more atoms at different elements
* Atoms combined in simple numerical ratios to form compounds
35
J.J. Thompson
- Plum Pudding/Raisin bread
| - e- in (+) framework
36
Ernest Rutherford
- discoverer of proton (Nuclear - Gold foil/a-scattering experiment)
- atom is mostly empty; (+) particles in nucleus
37
Neil Bohr
- Planetary
| - mostly used
38
Erwin Schrodinger
- Quantum/Mechanical/e- cloud
| - Modern atomic Model; estimates the probability of finding an e- in certain position (i.e. at e-cloud/ orbital)
39
Proton
- (+) ion
• Atomic number (basis of electronic configuration)
• Ernest Rutherford
40
Electrons
- (-) ion
• negligible weight 1,836x lighter that p+
- J.J. Thompson
- R.A. Millikan
41
J.J. Thompson
• Cathode ray tube: e- m/2 ratio
42
R.A. Millikan
Oil drop experiment: measure accurate charge and mass of e-
43
Neutrons
- no charge
• Atomic mass (Nucleon) = p+ + n0
• James Chadwick
44
Eugene Gold Stein
discovered anode rays
45
Electrochemistry
– particle separation based on e-
| • Ex: Capillary electrophoresis - separation of compounds based on electrophoretic mobility
46
ANODE (charge)
(+) electrode
47
ANODE (undergoes:)
Oxidation
48
CATHODE (charge:)
(-) electrode
49
CATHODE (undergoes:)
Reduction
50
RED CAT ELECT IN
• REDuction happens in CAThode where ELECTrons get IN
51
VILEORA
• Valence Increase, Loses e-, undergoes Oxidation, Reducing agent
52
VDGEROA
• Valence Decrease, Gains e-, undergoes Reduction, Oxidizing agent (KMnO4-, Na2Cr2O7)
53
Isotopes
* same p+/atomic number/ element
| * differ in atomic mass
54
Non-isotopes
19F, 127I, 31P, etc.
55
Main isotopes
+1: 1H, 12C, 14N, 32S, 35Cl ; +2: 16O, 79Br
56
Isobars
* same atomic mass
| * differ in elements
57
Isomers
* same molecular formula
| * differ in structure
58
Molecule
– aggregate of 2 or more atoms in definite arrangement held together by chemical bonds
59
Ions
– with net (+) or (-) charge
60
Empirical formula
– simplest whole number ratio (might be same with MF). Ex: CH2O vs. C6H12O6
61
Intermolecular FA/ Van der Waals/ Electrostatic
* Between molecule; weak and short-lived
| * Created by “molecule’s polarizability”; exerted when 2 uncharged atoms (n0) approach very closely
62
H-bonding
- Strongest IFA
| - H + S, O, N, X (electronegative atoms)
63
Keesom orientation (D-D)
- > (next to H-bonding)
| - Water-water
64
Debye Induction (D-ID)
- > (3rd)
| - Water-benzene
65
London Dispersion (ID-ID)
- Weakest IFA
| - Aromatics (Benzene-Benzene)
66
Dipole (D)
- Polar
67
Induced Dipole (ID)
- Nonpolar
68
Intramolecular FA
- Within molecule
69
Covalent
- sharing of e-
| - Nonmetal + Nonmetal (Glycosidic & Peptide bond)
70
Ionic
- Transfer of e-
| - Metal + Nonmetal (NaCl)
71
Glycosidic
- ether bond S─O─S
72
Peptide bond
– amide bond AA─peptide─AA bond
73
Covalent Bonding (Lone pair)
Pair of valence electrons that are not shared with another atom in covalent bond
74
Valence shell electron pair repulsion (VSEPR) theory
Predicts the geometry of the molecule as well as any bonded and unbonded electron pair
75
Linear (180˚)
- CO2
| - Alkynes (Sp)
76
Tetrahedral/bent (109.5˚)
- CCl4 , H2O Alkanes (Sp3)
| * 2 bonded pair, 2 unbonded pair
77
Trigonal bipyramid
- PF5
78
Octahedral
- SF5
79
Valence bond theory
States that bonds are formed by sharing of electron from overlapping atomic orbitals (covalent)
80
s = spherical
(sigma bond – stronger bond formed; headways overlap)
81
p = dumbbell
(pi bond – weaker; sideways overlap)
82
Molecular orbital theory
• States that bonds are formed from interaction of atomic orbitals from molecular orbitals
83
Bonding
- lower energy (stable)
84
Antibonding
- higher energy (unstable)
85
Synthesis/ Combination/ Direct Union
A + B → AB
86
Decomposition/ Analysis
AB → A + B
| - e.g. Complete & Incomplete combustion
87
Single Displacement
AB + X → AX + B
88
Double Displacement/ Metathesis/ Exchange
AB + CD → AC + BD
| - e.g. Neutralization, Precipitation
89
Metals
Li > K > Ba > Ca > Na > Mg > Al > Mn > Zn > Cr > Fe > Cd > Co > Ni > Sn > Pb > H2 > Cu > Ag > Hg > Pt > Au
90
Nonmetals (bases on electronegativity)
F > Cl > Br > I
91
Covalent compounds
* CO: Carbon monoxide
* SiO2: Silicon dioxide
* N2O: Dinitrogen monoxide
* CCl4: Carbon tetrachloride
92
Ionic compounds
Ex: Pb(NO3)4
• Classical: Plumbic nitrate
• Stock: Lead(IV) nitrate
93
Monovalent
```
+1 = Group 1 (H, Li, Na, K ׀ Ag)
+2 = Group 2 (Be, Mg, Ca, Sr, Ba ׀ Zn, Cd)
-2 = Group 6A (Oxide, Sulfide)
-1 = Group 7A (Fluoride, Chloride, Bromide, Iodide)
```
94
+1 = Group 1
(H, Li, Na, K ׀ Ag)
95
+2 = Group 2
(Be, Mg, Ca, Sr, Ba ׀ Zn, Cd)
96
-2 = Group 6A
(Oxide, Sulfide)
97
-1 = Group 7A
(Fluoride, Chloride, Bromide, Iodide)
98
Multivalent (with variable charges)
+1, +2 = Hg, Cu
+1, +3 = Au
+2, +3 = Fe, Co, Ni
+3, +5 = Bi, Sb
99
+1, +2
= Hg, Cu
100
+1, +3
= Au
101
+2, +3
= Fe, Co, Ni
102
+3, +5
= Bi, Sb
103
ClO-
- Hypochlorite
| - Hypochlorous acid (HClO)
104
ClO2-
- Chlorite
| - Chlorous acid (HClO2)
105
ClO3-
- Chlorate
| - Chloric acid (HClO3)
106
ClO4-
- Perchlorate
| - Perchloric acid (HClO4)
107
NO2-
- Nitrite
| - Nitrous acid (HNO2)
108
NO3-
- Nitrate
| - Nitric acid (HNO3)
109
SO3^2-
- Sulfite
| - Sulfite Sulfurous acid (H2SO3)
110
SO4^2-
- Sulfate
| - Sulfuric acid (H2SO4)
111
PO4^3-
- Phosphate
| - Phosphoric acid (H3PO4)
112
HCO3-
Bicarbonate (Hydrogen carbonate)
113
HSO3-
Bisulfite
114
HSO4-
Bisulfate
115
HPO4^-2
Biphosphate
116
H2PO4^-1
Dihydrogen phosphate
117
Aufbau Principle
* Atoms may be built by progressive filling of energy of main energy sub level (i.e., levels of lower energy levels are occupied first)
* s=2, p=6, d=10, f=14
118
Principal Quantum Number (n = 1 to 7)
* main energy level; size of orbital (electron cloud), distance of e- from nucleus
* Ex. O2 = 1s2. 2s2. 2p4 (n=2)
119
Azimuthal/ Angular Momentum (ℓ = 0 to 3)
* Angular momentum & shape of orbital; subshell
| * Ex. O2 = ℓ = 1
120
ℓ = 0 ─ s :
sharp (spherical shape)
121
ℓ = 1 ─ p :
principle (dumbbell shape)
122
ℓ = 2 ─ d :
diffuse (clover leaf)
123
ℓ = 3 ─ f :
fundamental
124
Magnetic Quantum Number (mℓ = -ℓ, 0, +ℓ)
* Orientation of orbital in space
| * Ex. O2 = mℓ = -1, 0, +1
125
Magnetic Spin (ms = + ½ , - ½ )
```
• Magnetic moment/ Rotation
Spin
↑ = Incomplete; clockwise + ½
↑↓ = Complete; counterclockwise = - ½
• Ex. Oxygen = ms= + ½
```
126
Diagmagnetism
– no unpaired e-
127
Paramagnetism
– at least 1 unpaired e-
128
• No 2 e- will have same set of quantum number (“exclusive”
Pauli’s exclusion theory
129
• Impossible to predict/ accurately determine the particle’s velocity (position & momentum)
Heisenberg’s uncertainty theory
130
* Orbitals are filled up singly before pairing up
| * Most stable arrangement of e- in subshells is the one with greatest no. of parallel spins.
Hund’s rule
131
P₁𝑽₁ = 𝑷₂𝑽₂ 𝑜𝑟 𝑷 ∝
1/𝑣
• Temperature (in K)
Boyle's/Mariotte
132
𝑻₁𝑽₁ = 𝑻₂𝑽₂
𝑜𝑟 𝑽 ∝ �
• Pressure (in atm)
Charles
133
𝑷₁𝑻₁ = 𝑷₂𝑻₂
𝑜𝑟 𝑷 ∝ 𝑻
• Volume (in L)
Gay-Lussac's
134
```
𝑷₁𝑽₁
𝑻₁
=
𝑷₂𝑽₂
𝑻₂
```
Combined
135
𝑷𝑽 = 𝒏𝑹𝑻
Ideal
136
• Equal volumes of different gases have same no. of moles at
STP
𝑽₁𝒏₁ = 𝑽₂𝒏₂
𝑜𝑟 𝑽 ∝ 𝒏 𝑜𝑟 𝑽/𝒏 = k
Avogadro's
137
• Total pressure in a mixture (non-interacting gases) is equal to the sum of the partial pressures of each gas. 𝑃𝑡 = 𝑃1 + 𝑃2 + 𝑃3
Dalton’s Law of Partial Pressures
138
• Rate of effusion (diffusion) and speed gas are inversely proportional to the square root of their density providing the temperature and pressure are same for 2 gases
Graham's
139
Rate at which 2 gases mix
Diffusion
140
Rate at which gas escapes through a pinhole vaccume
Effusion
141
• Diffusion rate (flux) of liquid or gas is directly proportional to the concentration gradient (ftom high concentration to low
concentration)
Fick’s 1st Law
142
𝑷𝒓𝒆𝒔𝒔𝒖𝒓𝒆 ∝ 𝑺𝒐𝒍𝒖𝒃𝒊𝒍𝒊𝒕𝒚
| • Decrease temperature, Increase Pressure (i.e., sealed container), more CO2 is dissolved in water.
Henry’s Law of Gas Solubility
143
Solute + Solvent
Solution
144
• Study of energy conversion/transformation in the universe
Thermodynamics
145
Allows exchange of energy and matter
Open System
146
Allows exchange of energy but not matter
Closed System
147
Does not allow exchange of both energy and matte
Isolated System "Adiabatic Walls"
148
* Independent (depends only on initial & final states of system)
* Ex. Enthalpy (H), Internal energy (U), Gibb’s Free Energy (G), Entropy (S)
State Function
149
* Dependent
| * Work and Heat
Non-state Function
150
• If two systems are in thermal equilibrium respectively with a third system, they must be in thermal equilibrium with each
other
Zeroth Law
151
• Energy is neither created nor destroyed but can be transformed from one form to another
1ST LAW: Law of conservation of Energy
152
∆H is independent of reaction/steps that occurred (only the initial and final steps is the basis)
Hess' Law
153
∆H = (+) → heat is absorbed;
COLD (endothermic)
154
∆H = (-) → heat is released;
HOT (exothermic)
155
* No way but UP
| * For an isolated system, Total entropy can never decrease over time
2ND LAW: Law of Entropy
156
Measure of system’s thermal energy per unit temperature; degree of disorderliness or randomness
Entropy (∆S)
157
* If an object reaches absolute zero temperature (0 K = -273.15 = -459.67 °)
* Entropy of perfect, solid, crystalline substance is zero at absolute 0 temperature
3RD LAW
158
* Thermodynamic state function that combines enthalpy and entropy
* ∆G = ∆H ‒ T∆S
Gibb’s free energy (∆G)
159
• Study of reaction rates and reaction mechanism
Chemical Kinetics
160
* Change in concentration of a reactant or product concentration with time
* aA + bB → cC + dD
Reaction Rate (M/s)
161
• Expresses relationship of the rate of reaction to the rate constant (K) and concentration of reactants raised to some
power
• aA + bB → cC + dD
Rate Law
162
• rate of chemical reaction is proportional to the number of collisions per time
Collision Theory
163
• (Formation of Intermediate Complex) - rate depends on Ea required to form intermediate state (where new bonds are
formed and old bonds are broken)
Transition Theory
164
= ↑ reactivity ↑ reaction rate (faster)
Nature of Reactants
165
= ↑ concentration ↑ reaction rate
Concentration of Reactants (except Zero order)
166
= ↑ reaction rate
Catalyst
167
= ↑ SA ↓ particle size ↑ reaction rate ↓ reaction time
Surface Area
168
```
= ↑ Temp ↑ KE ↑ mobility of molecules ↑ collision ↑ reaction rate;
Arrhenius Equation (T, Ea, RR)
```
Temperature
