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CHEM 2A > CHEM 2A: CH 1-10 > Flashcards

CHEM 2A: CH 1-10 Flashcards

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1
Q

Combination Reactions

A

2 Chemical species combining to become one

A+B —-> AB

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2
Q

Gases Behave Most Ideally at ____

A

low pressure and high temperature

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3
Q

Boyle’s Law

A

Relationship b/w pressure and volume

P1V1 = P2V2

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4
Q

Combined Gas Law

A

P1V1/T1 = P2V2/T2

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5
Q

PV = nRT

A 
P = atm
V = L
n = moles
R = (L)(atm)/(k)(mol)
T = Kelvin
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6
Q

Bond Order

A

1/2 number of valence electrons in bonding orbitals - number of electrons in anti-bonding orbitals

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7
Q

Molecular Orbital Theory (MO Theory)

A

Molecular orbitals are a combination of atomic orbitals, delocalized around the molecule

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8
Q

pi-Bond

A

P-P interaction, perpendicular to the internuclear axis

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9
Q

Sigma-Bond

A

Where the orbitals overlap in the internuclear axis

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10
Q

Hybrid Orbitals

A

Are localized (found on a specific point on molecules)

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11
Q

Hybridization

A

Atomic orbitals (s,p,d,f) mix together to give hybrid orbitals

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12
Q

Ground State

A

Lowest energy state

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13
Q

Orbital

A

3-d probability density predicting the most likely position where electrons cam be found

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14
Q

Decomposition Reactions

A

A single reactant compound breaks down into two or more product substances
AB—-> A + B

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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

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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
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17
Q

Reversibility

A

Indicated by the double arrow

————–>

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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

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19
Q

Single Replacement Reactions

A

One element replaces one of the ions in a compound

A + BX —–>AX + B

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20
Q

Neutralization Reactions

A

Reaction b/w acid and base (forms water and a salt)

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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)
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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
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23
Q

Chemical Properties/Changes

A

Properties: Flammability
Changes: Burns, Rust

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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
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25
Q

Percent Composition

A

%A = Parts A/Total Parts X 100

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26
Q

Isotopes

A

Two or more atoms of the same element that have different masses (different number of neutrons)

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27
Q

Dalton Atomic Theory

A

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

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28
Q

Law of Definite (or constant) Composition

A

Compound always made up in same proportion by mass

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29
Q

Significant Figures

A

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

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30
Q

Metric Prefixes

A 
Kilo- 1,000
Hecta- 100
Deca- 10
Meter- 1
Deci- 0.1
Centi- 0.01
Mili- 0.001
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31
Q

Ionic Bonds

A

Chemical bonds arising from the attractive forces b/w oppositely changed ions (interaction b/w cation and anion)

  • Strong
  • Solid @ room temperature
  • Crystal Lattice
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32
Q

Covalent Bonds

A

Chemical bond b/w 2 or more atoms that share a pair of electrons

  • Permanent until broken by chemical change
  • nonmetals interacting w/nonmetals
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33
Q

Cations VS. Anions

A

Cations- positive ions

Anions- negative ions

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34
Q

Temperature Conversions (F, C, K)

A 
TC= (5/9)(TF-32)
OK= -273.15 degrees C
TK= TC + 273
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35
Q

Limiting Reactants

A

Reactant that is totally consumed first, determines maximum yield possible

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36
Q

Solids

A
  • Stays same shape
  • Particles locked into place
  • Little free space b/w
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37
Q

Liquids

A
  • Volume remains constant
  • Shape changes
  • Particles move/slide past each other
  • Little free space b/w
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38
Q

Gases

A
  • Volume is variable
  • Move freely @ high speeds
  • Well seperated
  • Assumes shape of container
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39
Q

Particulate Matter

A

Smaller than microscopic atoms, H2O molecules, protein molecules

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40
Q

Microscopic Matter

A

Microscope, animal cells, plant cells

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41
Q

Macroscopic Matter

A

Matter= anything that has mass and takes up space

see w/eyes

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42
Q

Metric Units of Mass, Length, Volume

A 
Mass= Kg
Length= m
Volume= m^3
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43
Q

Polar VS. Non-polar Covalent Bonds

A

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

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44
Q

Ground VS. Excited State

A

Ground- the energy level an electron normally occupies

Excited - Release energy to go back to ground state

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45
Q

Electron

A
  • Negative Charge
  • 1/1837 mass H
  • e-
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46
Q

STP

A

273 K

1 atm

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47
Q

Percent Yield

A

Actual Yield/Ideal Yield X 100

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48
Q

Determining Molecular Formula

A

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

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49
Q

Determining Empirical Formula

A

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
Q

Valence Bond Theory

A

Number of bonds formed is exactly equal to the number of unpaired electrons found in the valence shell

51
Q

Bonding Scheme

A

Describes the actual orbitals used to make a bond

52
Q

Sigma Bond

A

1st bond formed b/w 2 atoms (always on the internuclear axis)

53
Q

Covalent (Molecular) Compound

A

Non-metal + non-metal bonding

  • lower boiling point
  • lower melting point
  • malleable
54
Q

Polar Bond

A

Bond b.w atoms where the 2 electrons are not shared equally (H2O)

55
Q

Formal Charge

A

number of free valence - number of electron it sees in molecules (lone pairs)

56
Q

Expanded Octet

A

Situation where the central atoms more than 8 electrons around it

57
Q

Double/Triple Bonds

A

Consist of one sigma (single) and any addition is pi bond

58
Q

Delocalize Electrons

A

Electrons that are free to move around

59
Q

Bond Energy

A

Energy require to break a bond

60
Q

Endothermic

A

When breaking bonds must add energy

61
Q

Exothermic

A

When bonds are formed energy is released

62
Q

Metallic Character

A
  • Decreases going crossed a period

- Increases going down a group

63
Q

Amphoteric

A

Can act as an acid or base (semi metal oxides, metalloids)

64
Q

Lewis Dot Diagram

A

Valence electrons are represented by dots and the element symbol

65
Q

Ionic Compound

A

Metal + Non-metal Bonding

  • Higher melting points
  • higher boiling points
  • Non-malleable (shatter/crumble)
66
Q

Electron Affinity

A

Amount of energy required to add and electrons to an atom in Gas Phase

67
Q

Ionization Energy

A

Amount of energy requires to remove and electron from an atom in gas phase

68
Q

Isoelectronic

A

Some number of electrons (larger atomic number will have a smaller raddii)

69
Q

Ionic Bond

A

Interaction b/w positive and negative ion

70
Q

Radia Nodes

A

Distances from the nucleus where electron density does not exist
- n-1-l

71
Q

Periodicity

A

Rises and falls in trends

72
Q

Atomic Radii

A

Across a period, size decrease

Down a group, size increases

73
Q

Paramagnetic

A

Where there is at least one unpaired electron

74
Q

Diamagnetic

A

All electrons are paired

75
Q

Angular Nodes

A

Areas with no electron density that’s associated with the shape of the orbital
- l

76
Q

Line Spectra (Atomic Spectra)

A

A series of lines that is displayed by an element

77
Q

Frequency

A

Number of cycles that pass through a given point of time

78
Q

Quantum Mechanical Model

A

How electrons exist in atoms

79
Q

l

A

Describes the shape of an orbital

80
Q

Charle’s Law

A

Relationship b/w volume and temperature

V1/T1 = V2/T2

81
Q

Graham’s Law of Effusion

A

the rate of A/the rate of B

82
Q

Emitted Photon

A

The one used to get to a lower energy (relaxation)

83
Q

Binding Energy (Threshold Energy)

A

How strongly the electrons is held to the surface (ionization energy)

84
Q

Gay- Lussac Law

A

P1/T1 = P2/T2

85
Q

Emission

A

Higher n to a lower n

86
Q

Quantum Numbers “ml”

A
  • Orbital magnetic number
  • Integer values from l—->l
  • Defines orientation
    “location in subshell”
87
Q

4f Orbitals

A 
n= 4
l= 3
ml= -2
ms= +1/2
88
Q

X-ray

A

the electromagnetic radiation that has the highest energy

89
Q

Angular Nodes

A

Areas w/no electron density that is associated w/ the shape of the orbital
= l

90
Q

Radial Nodes

A

Distance from the nucleus where the electron density does not exist

91
Q

Avogadro’s Law

A

Relationship b//w Number of moles of gas and volume

V1/n1=V2/n2

92
Q

Penetration

A

Orbitals which shield poorly will get pulled closer to the nucleus

93
Q

Quantum Numbers “n”

A

n= principle quantum number
- integer values (1 to infinity)
- energy shell (size)
“shell”

94
Q

Absorbed Photon

A

The one used to get to higher energy (absorption)

95
Q

Aufbau Princple

A

“buliding” up (fill lower energy before higher energy)

96
Q

Hund’s Rule (of Maximum of Multiplicity)

A

Degenerate orbitals singularly fill before pairing

97
Q

Pauli’s Exclusion Principle

A

All electrons in an toms must have a unique set of quantum numbers

98
Q

Shielding

A

How well an orbital can prevent the nucleus charge from interacting with electrons s>p>d>f

99
Q

Quantum Numbers “l”

A

Orbital Angular Momentum
- Integer values from 0 to n-1
- Defines type of orbital
“sub-shell” s,p,d,f

100
Q

Quantum Numbers “ms”

A

Spin magnetic #
- +1/2 or - 1/2
“Actual postion in an orbital”

101
Q

Mean Free Path

A

When mean free path decreases there is an increase on temperature. Average distance a gas particle can travel before a collision occurs.

102
Q

Deconstructive Interference

A

The waves cancel (out of phase)

103
Q

Constructive Interference

A

Waves add to give a larger wave (in phase)

104
Q

Wave-Particle Duality

A

Light can behave as a particle or a wave but not at the same time

105
Q

Spectator Ions

A

Species which do not participate in the RXN (the one’s that cancel out)

106
Q

Acid-Base Reactions (Neutralization)

A

Acid added to a base or base added to an acid

107
Q

Net Ionic Equation

A

Only species which actually participate in the reaction are included

108
Q

Total Ionic Equation

A

Compounds which are soluble are written interns of their ions

109
Q

Double Displacement

A

AB + CD —-> AD + CB

110
Q

Weak Acids

A

Are Weak Electrolytes

111
Q

Precipitation Reaction

A

Chemical RXN where a solid (precipitate) is formed when aqueous solutions are mixed

112
Q

Single Replacement/Displacement

A

AB + C—–> AC or CB

113
Q

Strong Acids

A

Are Strong Electrolytes

114
Q

Weak Electrolyte

A

Forms few ions

115
Q

Non-Electrylte

A

No ions formed

116
Q

Strong Electrolyte

A

Forms many ions

117
Q

Electrolyte

A

Compound which forms ions when dissolved in solvent

118
Q

Molecular Compounds

A

Do not dissolve into ions

119
Q

Solubility

A

Amount of solute that can be dissolved in a given amount of solvent

120
Q

Paschen Series

A

Is the emission spectrum of the hydrogen atom where the final energy level of all electrons in the n = 3 level

CHEM 2A (2 decks)

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