CHEM 2A: CH 1-10 Flashcards
1
Q
Combination Reactions
A
2 Chemical species combining to become one
A+B —-> AB
2
Q
Gases Behave Most Ideally at ____
A
low pressure and high temperature
3
Q
Boyle’s Law
A
Relationship b/w pressure and volume
P1V1 = P2V2
4
Q
Combined Gas Law
A
P1V1/T1 = P2V2/T2
5
Q
PV = nRT
A
P = atm V = L n = moles R = (L)(atm)/(k)(mol) T = Kelvin
6
Q
Bond Order
A
1/2 number of valence electrons in bonding orbitals - number of electrons in anti-bonding orbitals
7
Q
Molecular Orbital Theory (MO Theory)
A
Molecular orbitals are a combination of atomic orbitals, delocalized around the molecule
8
Q
pi-Bond
A
P-P interaction, perpendicular to the internuclear axis
9
Q
Sigma-Bond
A
Where the orbitals overlap in the internuclear axis
10
Q
Hybrid Orbitals
A
Are localized (found on a specific point on molecules)
11
Q
Hybridization
A
Atomic orbitals (s,p,d,f) mix together to give hybrid orbitals
12
Q
Ground State
A
Lowest energy state
13
Q
Orbital
A
3-d probability density predicting the most likely position where electrons cam be found
14
Q
Decomposition Reactions
A
A single reactant compound breaks down into two or more product substances
AB—-> A + B
15
Q
Law of Combining Volumes
A
When gases at the same temperature and pressure react, the reactant and product volumes in a ratio of small whole numbers
16
Q
Bohr Model of H Atom
A
- Electrons in orbitals of different energy around nucleaous
- Energy of electron are quantized on orbitals (not in between)
- Different energy levels hold different numbers of electrons
17
Q
Reversibility
A
Indicated by the double arrow
————–>
18
Q
Double Replacement Reactions
A
When solids of 2 compounds in aqueous solutions are mixed a positive ion from one species may combine with a negative ion from the other species to form a precipitate.
AX + BY ——> AY + BX
19
Q
Single Replacement Reactions
A
One element replaces one of the ions in a compound
A + BX —–>AX + B
20
Q
Neutralization Reactions
A
Reaction b/w acid and base (forms water and a salt)
21
Q
7 Strong Acids
A
HNO3 (Nitric Acid) H2SO4 (Sulfuric Acid) HCl (Hydrochloric Acid) HBr (Hydrobromic Acid) HI (Hydroiodic Acid) HClPO3 (Chloric Acid) HClO4 (Perchloric Acid)
22
Q
Scientific Notation
A
- 1 none-zero digit to the left of the decimal
- When decimal moves left, exponent increases
- When decimal moves right, exponent decreases
23
Q
Chemical Properties/Changes
A
Properties: Flammability
Changes: Burns, Rust
24
Q
Physical Properties/ Changes
A
- Observe with senses (height, length, width, smell, sight, size, color, boiling point)
- Changes: boil water, melt ice
- DOES NOT CHANGE SUBSTANCE
25
Percent Composition
%A = Parts A/Total Parts X 100
26
Isotopes
Two or more atoms of the same element that have different masses (different number of neutrons)
27
Dalton Atomic Theory
1) Each element is made up of individual atoms
2) Atoms are indivisible
3) All atoms of one element in every respect
4) Atoms of one element are different from atoms of any other element
5) Atoms combine to form compounds
28
Law of Definite (or constant) Composition
Compound always made up in same proportion by mass
29
Significant Figures
1) All non-zero digits are significant
2) All zeros b/w non-zero digits are significant
3) All zeros to the left of the 1st non-zero digit are NOT significant
4) Zeros to the right of the las non-zero digit when there is a decimal
5) Zeros to the right of the last non-zero digit are not significant if NO decimal
30
Metric Prefixes
```
Kilo- 1,000
Hecta- 100
Deca- 10
Meter- 1
Deci- 0.1
Centi- 0.01
Mili- 0.001
```
31
Ionic Bonds
Chemical bonds arising from the attractive forces b/w oppositely changed ions (interaction b/w cation and anion)
- Strong
- Solid @ room temperature
- Crystal Lattice
32
Covalent Bonds
Chemical bond b/w 2 or more atoms that share a pair of electrons
- Permanent until broken by chemical change
- nonmetals interacting w/nonmetals
33
Cations VS. Anions
Cations- positive ions
| Anions- negative ions
34
Temperature Conversions (F, C, K)
```
TC= (5/9)(TF-32)
OK= -273.15 degrees C
TK= TC + 273
```
35
Limiting Reactants
Reactant that is totally consumed first, determines maximum yield possible
36
Solids
- Stays same shape
- Particles locked into place
- Little free space b/w
37
Liquids
- Volume remains constant
- Shape changes
- Particles move/slide past each other
- Little free space b/w
38
Gases
- Volume is variable
- Move freely @ high speeds
- Well seperated
- Assumes shape of container
39
Particulate Matter
Smaller than microscopic atoms, H2O molecules, protein molecules
40
Microscopic Matter
Microscope, animal cells, plant cells
41
Macroscopic Matter
Matter= anything that has mass and takes up space
| see w/eyes
42
Metric Units of Mass, Length, Volume
```
Mass= Kg
Length= m
Volume= m^3
```
43
Polar VS. Non-polar Covalent Bonds
Polar Bonds- 2 atoms form a molecule using a covalent bond, shape electrons dipole-dipole, unequal share ionic charge
Non-Polar Bonds- equal share of electrons, no ionic charge
44
Ground VS. Excited State
Ground- the energy level an electron normally occupies
| Excited - Release energy to go back to ground state
45
Electron
- Negative Charge
- 1/1837 mass H
- e-
46
STP
273 K
| 1 atm
47
Percent Yield
Actual Yield/Ideal Yield X 100
48
Determining Molecular Formula
1) Determine Empirical Formula
2) Calculate molar mass of the Empirical Formula Unit
3) Determine molar mass of the compound
4) Divide MM of compound by MM E.F. (n)
5) Multiply by empirical formula
49
Determining Empirical Formula
1) Determine % composition by mass of each element in sample
2) Convert masses into moles
3) Determine ratio of the moles of atoms
4) Express the moles of atoms as the smallest possible ratio of integers
50
Valence Bond Theory
Number of bonds formed is exactly equal to the number of unpaired electrons found in the valence shell
51
Bonding Scheme
Describes the actual orbitals used to make a bond
52
Sigma Bond
1st bond formed b/w 2 atoms (always on the internuclear axis)
53
Covalent (Molecular) Compound
Non-metal + non-metal bonding
- lower boiling point
- lower melting point
- malleable
54
Polar Bond
Bond b.w atoms where the 2 electrons are not shared equally (H2O)
55
Formal Charge
number of free valence - number of electron it sees in molecules (lone pairs)
56
Expanded Octet
Situation where the central atoms more than 8 electrons around it
57
Double/Triple Bonds
Consist of one sigma (single) and any addition is pi bond
58
Delocalize Electrons
Electrons that are free to move around
59
Bond Energy
Energy require to break a bond
60
Endothermic
When breaking bonds must add energy
61
Exothermic
When bonds are formed energy is released
62
Metallic Character
- Decreases going crossed a period
| - Increases going down a group
63
Amphoteric
Can act as an acid or base (semi metal oxides, metalloids)
64
Lewis Dot Diagram
Valence electrons are represented by dots and the element symbol
65
Ionic Compound
Metal + Non-metal Bonding
- Higher melting points
- higher boiling points
- Non-malleable (shatter/crumble)
66
Electron Affinity
Amount of energy required to add and electrons to an atom in Gas Phase
67
Ionization Energy
Amount of energy requires to remove and electron from an atom in gas phase
68
Isoelectronic
Some number of electrons (larger atomic number will have a smaller raddii)
69
Ionic Bond
Interaction b/w positive and negative ion
70
Radia Nodes
Distances from the nucleus where electron density does not exist
- n-1-l
71
Periodicity
Rises and falls in trends
72
Atomic Radii
Across a period, size decrease
| Down a group, size increases
73
Paramagnetic
Where there is at least one unpaired electron
74
Diamagnetic
All electrons are paired
75
Angular Nodes
Areas with no electron density that's associated with the shape of the orbital
- l
76
Line Spectra (Atomic Spectra)
A series of lines that is displayed by an element
77
Frequency
Number of cycles that pass through a given point of time
78
Quantum Mechanical Model
How electrons exist in atoms
79
l
Describes the shape of an orbital
80
Charle's Law
Relationship b/w volume and temperature
| V1/T1 = V2/T2
81
Graham's Law of Effusion
the rate of A/the rate of B
82
Emitted Photon
The one used to get to a lower energy (relaxation)
83
Binding Energy (Threshold Energy)
How strongly the electrons is held to the surface (ionization energy)
84
Gay- Lussac Law
P1/T1 = P2/T2
85
Emission
Higher n to a lower n
86
Quantum Numbers "ml"
- Orbital magnetic number
- Integer values from l---->l
- Defines orientation
"location in subshell"
87
4f Orbitals
```
n= 4
l= 3
ml= -2
ms= +1/2
```
88
X-ray
the electromagnetic radiation that has the highest energy
89
Angular Nodes
Areas w/no electron density that is associated w/ the shape of the orbital
= l
90
Radial Nodes
Distance from the nucleus where the electron density does not exist
91
Avogadro's Law
Relationship b//w Number of moles of gas and volume
| V1/n1=V2/n2
92
Penetration
Orbitals which shield poorly will get pulled closer to the nucleus
93
Quantum Numbers "n"
n= principle quantum number
- integer values (1 to infinity)
- energy shell (size)
"shell"
94
Absorbed Photon
The one used to get to higher energy (absorption)
95
Aufbau Princple
"buliding" up (fill lower energy before higher energy)
96
Hund's Rule (of Maximum of Multiplicity)
Degenerate orbitals singularly fill before pairing
97
Pauli's Exclusion Principle
All electrons in an toms must have a unique set of quantum numbers
98
Shielding
How well an orbital can prevent the nucleus charge from interacting with electrons s>p>d>f
99
Quantum Numbers "l"
Orbital Angular Momentum
- Integer values from 0 to n-1
- Defines type of orbital
"sub-shell" s,p,d,f
100
Quantum Numbers "ms"
Spin magnetic #
- +1/2 or - 1/2
"Actual postion in an orbital"
101
Mean Free Path
When mean free path decreases there is an increase on temperature. Average distance a gas particle can travel before a collision occurs.
102
Deconstructive Interference
The waves cancel (out of phase)
103
Constructive Interference
Waves add to give a larger wave (in phase)
104
Wave-Particle Duality
Light can behave as a particle or a wave but not at the same time
105
Spectator Ions
Species which do not participate in the RXN (the one's that cancel out)
106
Acid-Base Reactions (Neutralization)
Acid added to a base or base added to an acid
107
Net Ionic Equation
Only species which actually participate in the reaction are included
108
Total Ionic Equation
Compounds which are soluble are written interns of their ions
109
Double Displacement
AB + CD ----> AD + CB
110
Weak Acids
Are Weak Electrolytes
111
Precipitation Reaction
Chemical RXN where a solid (precipitate) is formed when aqueous solutions are mixed
112
Single Replacement/Displacement
AB + C-----> AC or CB
113
Strong Acids
Are Strong Electrolytes
114
Weak Electrolyte
Forms few ions
115
Non-Electrylte
No ions formed
116
Strong Electrolyte
Forms many ions
117
Electrolyte
Compound which forms ions when dissolved in solvent
118
Molecular Compounds
Do not dissolve into ions
119
Solubility
Amount of solute that can be dissolved in a given amount of solvent
120
Paschen Series
Is the emission spectrum of the hydrogen atom where the final energy level of all electrons in the n = 3 level
