Final Exam Flashcards
(235 cards)
1
Q
Allotropes
A
Allotropes: different forms of same element
2
Q
Pure substance:
A
constant & uniform composition
3
Q
Pure substance: ELEMENT vs. COMPOUND
A
Element: molecules contain only 1 type of atom [O2]
Compounds: combining different atoms [H2O] –> can be decomposed into 2+ atoms
4
Q
Mixture:
A
2+ types molecules can be separated by physical changes (ie. evaporation)
5
Q
Mixture: Homogenous vs Heterogeneous
A
Homogenous mixture: uniform composition - solutions & air (“solution” ie. air)
Heterogeneous mixture: composition varies from point to point
6
Q
Physical properties:
A
observed without changing substance into another, reversible
(mass, volume, density, boiling pt, solubility, color)
7
Q
Physical properties: intensive vs extensive
A
Intensive: independent amount substance (boiling pt, density, color)
Extensive: depend amount substance (weight, mass, length)
8
Q
Chemical properties:
A
undergoes change chemical composition
(flammability, corrosiveness, reactivity with acid)
Key words: reacting, changing, burning
9
Q
Changes in matter: physical vs chemical
A
Physical change: don’t change composition of substance and no new substance is formed (wax melts, magnetizing solids)
Chemical change: result in formation of new substance with different chemical properties (combustion, oxidation)
10
Q
Significant figure rules
A
All non-zero digits = significant
Zeros:
Left = not significant
Middle = significant
Right = significant after decimal point
11
Q
Rounding numbers rules
A
Adding/subtracting: same # decimal places as the # with least decimal places
Multiplying/dividing: same # of digits as # with least sig figs
If digit dropped < 5 (round down) if > 5 (round up)
If digit dropped = 5 (round up or down whichever yields an even value for the retained digit)
12
Q
Accuracy
Precision
Exact numbers
A
Accuracy: variation between experimental and accepted value
Precision: how similar results are when repeated in the same manner
Exact numbers: counting, definition, unit conversion (infinite # sig figs, no uncertainty)
13
Q
Density
A
desnity = mass / volume
[g/ml = g/cm^3]
Used to identify unknown substance
Determine density of irregular object: use volume water displaced in beaker
14
Q
Atom
A
Atom: smallest unit of an element that can participate in chemical change (indivisible)
15
Q
Element vs molecule
A
Element: 1 type of atom, in which mass is a characteristic feature
Molecule: 2+ atoms joined by chemical bonds
16
Q
Dalton’s atomic theory:
A
WRONG, but laid foundation for future work
- All matter is made up of tiny particles called atoms (indivisible & indestructible)
- Atoms given element are identical in size, mass, chemical properties
- Atoms combine to form compounds in whole number ratios
- Atoms of element cannot change into atoms another element (only rearrange)
17
Q
Cathode ray tube, J.J. Thomson
A
DISCOVERED: electrons negative, charge to mass ratio electron
Plum pudding model atom
18
Q
Oil drop experiment, Millikan:
A
DISCOVERED charge of an electron
19
Q
Alpha-ray scattering, Rutherford
A
DISCOVERED nucleus in atoms (disprove plum pudding)
Nucleus has protons & neutrons (which are much much heavier than electrons) → nucleus accounts for most of an atom’s mass but very little of its size
20
Q
Rutherford’s model of atom:
- 3 main points
- limitations
A
- All + charge & mass concentrated inside nucleus (tiny region)
- Negatively charged particles revolve around nucleus in circular path
- Electrostatic force attraction between proton & electrons holds atom together
LIMITATIONS: failed to explain
- Stability of an atom
- Electronic structure of atom (how electrons arranged inside nucleus)
21
Q
Atomic number (Z):
“Nuclear charge”
A
protons in nucleus (found in periodic table)
Neutral atom: electrons = protons = Z
22
Q
Mass number (A):
A
protons + neutrons
23
Q
Nuclear symbol:
A
represents nucleus of an isotope
Atomic number (Z): protons
Mass number (A): protons + neutrons
24
Q
Empirical formula:
A
molecular formula expressed in lowest whole number ratio
25
Mole:
number of atoms/molecules in a bulk sample of matter
26
Avogadro’s number
Number of particles per mole
27
Isotopes:
atoms with same atomic number (Z) but different mass number (A)
carbon 12, 13, 14
28
Percent abundance of isotope
29
Atomics mass:
weighted average of isotopic masses (mass spectroscopy)
30
Electromagnetic (EM) radiation:
oscillating electric & magnetic field perpendicular to each other & direction propagation → quantized (units: photons)
31
Wavelength (λ):
distance between 2 consecutive peaks/troughs
[m, nm]
32
Frequency (v)
cycles pass through given point / second
[Hz = 1/sec]
33
Amplitude
height of peak, corresponds to brightness/intensity
34
All types of EM radiation travel at what speed?
Speed of light
3 x 10^8 m/s
35
Electromagnetic spectrum:
microwave → infrared → visible → ultraviolet → x-ray → gamma ray
36
Wavelength range of visible light:
(LOW frequency) ROYGBIV (HIGH frequency)
37
Photoelectric effect:
light wave is particulate in nature, consisting of small packets of energy called photon
CORRECT
- Electrons ejected from metal when light has frequency greater than threshold (>0)
- Energy proportional to number of photons
38
Wave theory (classical physics):
WRONG
- Energy light should correspond to intensity, nothing to do with frequency
- Kinetic energy of electron shouldn't change with frequency, change with intensity of EM radiation:
- KE electrons should increase linearly with intensity of light
- # electrons should be independent of intensity & brightness incident light
39
Photoelectric effect: problems
40
Continuous vs discontinous spectrum
Continuous spectrum: contains all wavelengths of visible light
Discontinuous spectrum: missing/discontinuous wavelength (line spectra)
41
Line spectra vs absorption spectra
Line spectra (emission): atoms release energy
Absorption spectrum: atoms absorb energy
42
Bohr’s model of atom:
combines classic & quantum physics:
- Stationary states: electrons move around nucleus in ORBITS
- Quantization of angular momentum: electron revolves around nucleus only in specific orbits in which angular momentum of electron is an integral multiple of h/2π
- Energy levels: unequal spacing (radius decreases by n^2), closer electron is to nucleus = smaller energy
- When electrons jump lower → high orbit (absorbs energy), higher → lower (emit energy)
43
Bohr’s model of atom: equation for amount of energy abosrbed/emitted when electron jumps between energy levels
When electrons jump lower → high orbit (absorbs energy), higher → lower (emit energy)
44
Bohr’s model of atom:
Radius of n’th orbit of MONO electronic atoms (1 electron):
45
Bohr’s model of atom:
Quantization of angular momentum
46
What does Ryberg's equation tell?
find wavelength of photon resulting from an electron jumping between energy levels
47
What does Bohr’s energy equation say?
When electron jumps between energy levels, energy of photon emitted = energy transition
48
Spectral line: Lyman series
electron jumps from higher energy level → ground state (n = 1)
UV light
49
Spectral line: BALMER series
electron jumps from higher energy level → n=2
Visual light
50
Spectral line: Paschen, Bracker, Pfund
Paschen series --> n = 3
Bracker series: --> n = 4
Pfund series: --> n = 5
INFRARED light
51
How do you compare the wavelength of the 1st line in balmer series to that of lyman series?
Wavelength of the 1st line in Balmer series > first line of Lyman series
As we move higher n the ∆E for next shells decreases
52
Heisenberg uncertainty principle (HUP)
impossible to know accurately & simultaneously both the position & momentum of moving particle
Rules out existence orbits with specific radius of electrons (instead use probabilities to express electron’s position)
53
Bohr’s explanation of spectral lines:
- Electron excited it jumps to higher energy level
- Higher energy level less stable --> jump back down to ground state (n = 1), releasing photon
- Photon appears as spectral line on emission spectra
54
Shrodinger wave equation
Shrodinger wave equation: describe electron wave in 3D
Square of wave function: probability of finding an electron at a particular point
Possible solutions wavefunction (orbitals) each has unique energy (E)
55
Energy of electron in hydrogen atom important points:
1. Only dependent on principle quantum number (n), does not depend on l or ml
2. Energy levels are quantized → can only have certain discrete energy values
56
Orbital:
3D space around the nucleus where the probability of finding an electron is max.
57
What do we need to define a particular orbital?
3 quantum numbers
n: shell, size & energy
l: subshell, shape of orbital
ml: orientation of orbital
58
Principle quantum number (n)
(n): size & energy of the orbital
positive integer n=1,2,3..
As value of n increases, size & energy of that orbital increases
max #orbitals with a given n value = n^2
(each orbital has 2 electrons)
59
Angular momentum quantum number
(l): shape of the orbital
for a given n, l is between 0 and n-1
(in other words n>l)
Orbitals with same l value are in same subshell
60
Hierarchy: shells, subshells, orbitals
shells → subshells → orbitals
61
Magnetic quantum number
For a particular subshell how many possible orbitals are in subshell?
(ml): orientation of the orbital
ml can be -l to +l (including 0)
(in other words l > ml)
For a particular subshell with defined l value, there are (2l+1) orbitals in subshell = possible values of ml
62
Electron spin quantum number
(ms): spin of the ELECTRON, says nothing about the ORBITAL
2 different orientations +1/2,-1/2 (↑, ↓)
Spin represents an intrinsic property of the electron, NOT property of an orbital (like the other 3 quantum #s)
63
Each orbital contains max how many electrons?
max 2 electrons
64
65
Ways to visualize orbitals: 3D space around nucleus where probability of finding electron is max
1. Electron dot plot: desnity = probability finding electron (higher closer to nucleus)
2. Probaility plot: square wavefunction vs radius
3. Radial probability distribution (RPD): probability of finding an electron in a spherical shell of thickness dr at a distance r from nucleus
66
Node:
Total nodes
Radial nodes
Angular nodes
region with zero electron probability
total #nodes = n - 1
radial nodes = n - 1 - l (spherical)
angular nodes = l (plane)
67
Degenerate orbitals:
same n → same amount of E
68
What does energy & orbital depend on in multielectron atom?
energy of an electron in a single atom can be determined solely by the principal quantum number (n = shell)
energy electron in multielectron atom depends on principle quantum number n (shell) & angular momentum quantum number l (subshell)
Different subshells can have different energies (s
69
Do subshells have the same amount of energy?
Different energies (s < p < d < f)
70
Shielding effect
Effective nuclear charge
Shielding effect: reduction of nuclear charge (Z) to the effective nuclear charge (Zeff) by other electrons in a multi-electron atom
Effective nuclear charge: pull exerted on an outer electron by the nucleus, taking into account electron-electron repulsion
Zeff = Z - S = protons - shielding electrons (between valence & nucleus)
Inner electrons shield outer electrons much more than electrons in the same shell shield each other
71
Penetration effect
Distance electrons inside a particular orbital is to nucleus
Closer electron is to nucleus, lower energy associated with orbital
72
Order penetration of orbitals
Order of orbital energy
Order penetration of orbitals: s>p>d>f
Order of orbital energy: s
73
Electron configuration
how electrons of an atom are filled into atomic orbitals
74
Aufbau principle
Pauli’s exclusion principle
Hund’s rule
Aufbau principle: fill electrons in lowest energy orbitals first
Pauli’s exclusion principle: in a given atom, 2 electrons cannot have the same set of 4 quantum numbers (n,l,ml,ms)
- Orbital hold max 2 electrons with opposite spins
Hund’s rule (degenerate orbitals): when filling electrons into orbitals of equal energy, fill each orbital with a single electron, maintaining parallel spins (up), before doubling up electrons in that orbital set (down)
75
Electron configuration for transition elements
loose s electrons before losing d electrons
76
Anomalous electron config: unexpected, exceptions (cost additional energy)
Completely filled or ½ filled subshell: more stable (electron shifts to a higher energy orbital)
Cr, Mo, Cu, Ag
77
As we go across a period
add proton to nucleus & electron to valence shell with each element
78
Periodic Properties: Atomic size/radius
Down group (↓) valence electrons are in larger orbitals
Across period (←): # protons increases while # shielding electrons remains same therefore pulled closer to nucleus
79
Isoelectronic series:
series of atoms/ions that have same number of electrons therefore same shielding
Most negative ion has largest radius, most positive ion has smallest (within isoelectronic set)
80
Ionization energy:
energy required to remove an electron from a neutral atom in its gaseous phase [kJ/mol]
IE1 < IE2 ... (every electron removed decreases shielding, decreases atomic radius, increases next IE)
81
What is a way of identifying number valence electrons in atom using IE?
Removing an electron from inner shell >> valence electron (much more energy)
identify #v.e by seeing where BIG PEAK in IE occurs
82
Electron affinity:
energy released when an electron is added to a valence shell of the atom
Halogens: high negative EA (more likely to gain electron)
Noble gasses: large positive EA (add electron requires energy, unlikely to gain electron)
83
Ionization energy (IE) exceptions
Add 1st v.e p-subshell (shielded by full s-subshell): easy to remove
1st set paired electrons formed in p-subshell (electron electron repulsion): 4th electron easy to remove
84
Why do atoms form chemical bonds?
- Decrease energy
- Increase stability
85
Ionic crystal lattice:
3D arrangements of cations & anions held together by electrostatic force of attraction
Ionic solids: high melting/boiling point, poor conductor electricity (solid), excellent conductors electricity (dissolved/melted
86
Molecules
smallest unit of covalent compound
87
Bond length
combined energy of both bonding atoms is min
distance which the lowest potential energy is achieved
88
Bond energy (BE):
energy required to break bond or energy released when bond is formed
89
Breaking chemical bonds vs forming chemical bonds, energy?
Break chemical bonds (energy must be added, ENDOTHERMIC)
Forming chemical bonds (release energy, EXOTHERMIC)
90
Bond order:
electrons shared between a pair of atoms
91
Ionic bond:
NM & M (transfer of electrons)
EN difference = +1.7
92
Covalent bond
NM & NM (sharing electrons)
EN difference = less 1.7
93
Non polar vs polar covalent bond
Non polar covalent bond: between same atoms - equal sharing electrons
Polar covalent bond: unequal sharing electrons - partial negative charge most EN
94
ELECTRONEGATIVITY:
tendency of atom to attract electrons towards itself
Greater electronegativity (∆EN) = greater polarity of bond
95
Naming: Ionic compounds
Basic:
Includes transition metal:
Include polyatomic ions:
Basic: M NM + ide
Includes transition metal: M (charge) NM + ide
Include polyatomic ions: M + polyatomic ion name
96
Naming: Binary covalent compounds ( 2 types of NM)
Prefix NM prefix NM + ide
Prefix: mono (1), di (2), tri (3), tetra (4), pentra (5), hexa (6), hepta (7), octa (8)
97
Lattice energy:
energy released when 1 mol of ionic crystal is formed from its cations & anions in their gasous phases (g)
Stronger F means higher lattice energy (to pull apart)
98
Lewis Dot Structure rules
- Atom lowest electronegativity in center (except for H cannot be placed as center)
- Add +1 v.e for negative charge, subtract -1 v.e for positive charge
99
lewis exceptions
100
Resonance structures
same arrangement of atoms but diff placement of electrons
Resonance hybrid: superposition/avg of resonance structures
101
What is wrong about how lewis structures depict electrons?
Lewis structures depict electrons as localized between given pair of atoms (bond) or an individual atom (lone pair)
In nature, electrons are delocalized: density spread over entire molecule
102
Bond order:
electron pairs shared between two atoms
103
Equivalent vs non-equivalent resonance structure
same distribution formal charges
Non-equivalent resonance structure: different distribution FC (Don’t equally contribute to resonance hybrid)
104
Formal charge (FC):
hypothetical charge an atom would have if we could redistribute the electrons in the bonds evenly between the atoms
FC= v.e free atom - v.e bonded atom
Sum formal charges on all atoms = overall charge molecule
105
Criteria for choosing greatest contribution to resonance hybrid:
Smaller FC are preferable
Same nonzero FC on adjacent atoms not preferable
A more negative FC should reside on more electronegative atom
106
Choose a likely identity for X in these structures:
Count total number of v.e in structure
Find number of v.e that X contributes to structure
v.e = group that atom belongs to
107
Drawbacks of lewis structure:
doesn’t explain shape/geometry of molecule (effect properties)
108
Valence Shell Electron Pair Repulsion Theory (VSEPR):
predict shape/geometry molecule from lewis dot structure
electron pairs are located as far apart from each other as possible → reduces repulsions between electron pairs → decrease potential energy molecule → increases stability
Bonds/lone pairs counted 1 valence shell electron pair (no matter if it’s a single or triple bond)
109
Bond strength:
Bond size:
Strength: Triple bond > Double bond > Single bond
Size: Triple bond < Double bond < Single bond
110
Electron pair geometry vs Molecular geometry
Electron geometry: electron pairs around central atom (includes lone pairs & bonding pairs)
Molecular geometry: atoms relative to central atom (excludes lone pairs)
111
VESPR chart
Nonpolar highlighted in blue (if all atoms are same)
112
Valence bond theory (VB):
Bonding atoms approach each other atomic orbitals overlap
Each atomic orbital that overlap has 1 electron opposite spin
After orbitals overlap, pair of electrons occupy the overlapped region
113
How to determine strength bond in VB theory?
greater orbital overlap (closer nuclei is to bonded electrons) = stronger bond
114
Sigma bonds (σ):
covalent bond - electron density internuclear axis
(single bond = σ bond)
Lobes point toward each other (end to end overlap), free rotation around sigma bond
115
Pi bond (π):
side by side overlap of 2 p orbitals - electron density above and below internuclear axis
(double bond = 1 + 1π )
π bond weaker than σ bond (sideways overlap less effective)
116
Hybridization:
orbital mixing to form hybrid orbitals (linear combination of atomic orbitals)
- Number hybrid orbitals formed = number of atomic orbitals combined
- Hybrid orbitals formed are equivalent in shape and energy
- Hybrid orbitals are more effective in forming bonds than unhybridized orbitals
- Hybrid orbitals orient themselves in 3D to max distance between then and min repulsions between electrons
Number of electron groups around central atom = number of hybrid orbitals
117
Hybridization:
triple, double, single bond
118
Sp hybridized orbitals on Be in BeCl2 molecule
1s orbital + 1p orbital
119
Sp2 hybridization in BH3 atom
1s orbital + 2p orbital
120
Sp^3 hybridization in CH4
1s orbital + 3p orbital
121
Does resonance influence hybridization?
No! arrangement of π bonds involves only the unhybridized orbitals
Hybridization involves only σ bonds, lone pairs of electrons, and single unpaired electrons
122
Polar molecule:
asymmetric distribution or different atoms
Dipole moment: measure polarity
Arrow pointing toward more electronegative
- Any molecule with lone pairs of electrons around the central atom is polar.
123
Nonpolar molecule
symmetric distribution, EN = 0, no lone pairs
124
Intermolecular forces:
attraction between molecules with partial charges, or between ions/molecules
125
Dipole dipole moment
between POLAR molecules
126
London dispersion forces
between ALL atoms due to random motion of electrons
-Stronger LDF in larger, heavier molecules
-Branched molecules have weaker LDF than straight chain
127
Hydrogen bond
strong dipole–dipole
H & N, O, F (covalent bond)
128
Molecular orbital theory (MO):
atomic orbitals in bonding atoms combine to form molecular orbitals (DELOCALIZED)
atomic orbitals combined = # of molecular orbitals created
Bonding MO lower energy than atomic orbitals b/c increased stability associated with bond formation
129
Valence bond theory vs. molecular orbital theory
130
Filling molecular orbitals:
MOs are filled in order of increasing energy (Aufbau principle)
An MO can hold max of 2 electrons with opposite spin (Pauli exclusion)
Degenerate orbitals are ½ filled with parallel spins before doubling up (Hund's rule)
131
Bonding orbital
electron density directly between nuclei
Placing electron in bonding orbital stabilizes molecule b/c between 2 nuclei
132
Antibonding orbital (*)
placing electron in non-bonding orbital destabilises
133
Degenerate orbitals:
Electron orbitals having the same energy levels
orbitals in 2p subshell: 2px, 2py, 2pz
134
Delocalized π bond
π bond extend over 2+ atoms
In lewis structure, this occurs when with resonance structures involving double & triple bonds
135
Molecular orbital diagram:
shows energy and number of electrons present in each MO, atomic orbitals from which each MO is formed
136
P atomic orbiatls --> MO orbitals formed
137
Bond order
If bond order > 0: stable
Higher bond order = stronger bond
138
Paramagnetic:
Diamagnetic:
Paramagnetic: has unpaired electrons
(MO theory can explain paramgnetic, lewis can't)
Diamagnetic: all electrons paired
139
What is avogadro's number?
number atoms or molecules / moled
140
Percent composition:
% by mass of each element in compound (defines identity)
141
Chemical formula
Empirical formula
Chemical formula: how many atoms of each element are in a compound
Empirical formula: simplest ratio of elements in a compound (divide by smallest number)
142
Empirical formula mass
avg atomic masses of all atoms in empirical formula
143
Limiting reactant:
completely consumed first, limits amount of product formed
ICE reaction table: initial, change, end
144
Theoretical, actual, percent yield
145
Solutions:
Solute:
Solvent:
Solutions: homogeneous mixture of 2+ substances
Solute: smaller amount
Solvent: larger amount (ie. water is universal solvent)
146
Ionic compounds in water:
Forces of attraction between solute & solvent particles dissolve
Not all ionic substances are soluble in water
Solubility depends on attraction between ions & water
147
Strong electrolytes
dissociate completely into ions (conduct electricity)
- Soluble ionic compounds (NH4Cl)
- Strong acids & bases (HCl)
148
Weak electrolytes:
dissociate into fewer number of ions (don’t conduct well)
- Weak acids & bases (HF)
- Insoluble ionic compounds (AgCl)
149
Non electrolytes
Non electrolytes: dissolve in water (polar) but don’t produce ions (don’t conduct)
- Covalent compounds (C12H22O11)
150
Molarity (M)
moles of solute in 1L of solution
151
How to prepare 250mL of a 0.200M aqueous solution of sucrose (aq):
152
Dillute vs concentrated
Dilute: small amount of solute dissolved
Concentrated: large amount of solute dissolved
153
Dillution
Dilution: adding additional solvent to a solution
moles doesn’t change during dilution
154
Chemical reaction types:
Precipitation reactions: (aq) + (aq) → (s)
Acid base reactions: acid + base → salt + water (l)
Redox reaction: change in oxidation number
Combustion reaction: CxHy + O2 → CO2 + H2O
155
Molecular equations
Complete ionic equations
Net ionic equation
Molecular equations: all reactants/products as if they were intact, undissociated compounds
Complete ionic equations: all soluble ionic substances dissociated into ions
Net ionic equation: eliminates spectator ions, shows only actual chemical change
156
Alkali metal cations:
Li, Na, K, Rb, Cs, Fr
157
Example: predict if a precipitate will form when aqueous solution of sodium nitrate (NaNO3) is combined with aqueous solution of potassium iodide (KI). If yes, write molecular, complete ionic, and net ionic equations.
158
Arrhenius definition of acid & base limited to aqueous solutions
Arrhenius acid: acid is a substance that produces H+ ions (protons) in water
Arrhenius base: substance that produces OH- ions in water
159
Bronsted-lowry definition acid & base
Acid: proton donor (loose H very easily b/c weak bonds → become weak bases)
Base: proton acceptor → have a lone pair (hold H very tightly b/c strong bonds → become weak acid)
160
Strong electrolytes:
dissociate completely in aq solution
Acids & bases can act as electrolytes (conduct electricity)
161
Electrolyte examples
strong acids
weak acids
strong bases
weak bases
Strong acid: HCl HBr Hi & as long as you have 2+ O than H
Strong base: B pt table
162
Neutralization reaction:
acid reacts with a base to form salt & water
acid (aq) + base (aq)→ salt (aq) + H2O (l)
Net ionic equation: (H+) + (OH-) → H2O
163
Titration:
titrant, analyte
analytical technique to determine concentration of unknown acid/base solution
Titrant: solution known concentration
Analyte: solution unknown concentration
164
Equivalence point:
moles acid = moles base
amount titrant added completely neutralizes analyte solution
Acid & base completely consumed and neither of them are in excess
At equivalence point:
- Strong acid neutralizes weak base: solution pH < 7
- Strong base neutralizes weak acid: solution pH > 7
- Strong acid neutralizes strong base: solution pH = 7
165
How is titration done?
166
How to determine equivalence point?
167
Redox reactions (oxidation reduction reaction):
transfer of electron(s) from one species to another
change in oxidation number (ON)
168
169
Oxidation number (ON):
Any lone element = 0
Assign most EN element first, ON correspond to group #
Hydrogen: +1 (NM), -1 (M)
Sum ON = charge
170
Important characteristics of redox reaction:
oxidation & reduction must occur together
171
Oxidation vs reduction
OIL RIG
Element that is OXIDIZED undergoes increase in ON (loss electrons)
Element that is REDUCED undergoes decrease in ON (gain electrons)
172
Cation displacement reaction
"Single replacement reaction"
metal with high reactivity in the activity series is added to a solution containing a cation with lower reactivity
AB (aq) + C (s) → A (s) + CB (aq)
predict a single replacement reaction will occur when a less reactive element can be replaced by a more reactive element in a compound
173
Atmospheric pressure:
pressure exerted by column if air from top of atmosphere to surface of earth (higher altitude = smaller atmospheric pressure)
Barometer device to measure atmospheric pressure → pressure indicated by height (mm) of mercury column
174
Gas pressure
Manometer device to measure pressure of gas inside container
175
Boyles law
P inversely proportional V
P1V1 = P2V2
176
Charles law
V proprtional to T (K)
177
Avogadro's law
V proptional to n
equal volumes of any ideal gas contain equal number of particles (moles)
178
Ideal gas law
PV = nRT
Temperature in K
179
Standard temp & pressure (STP):
0ºC, 1atm
180
Standard molar volume:
1 mole of any ideal gas at standard temp & pressure occupies 22.4L
181
Determine molar mass of gas using ideal gas law:
182
Dalton’s law of partial pressure
mixture of non-reacting gasses - total pressure is equal to the sum of the pressures that each gas would exert if it were alone
P = Pa + Pb...
183
Mole fraction
Value mole fraction between 0 and 1, sum mol fraction of all components must add up to 1
184
Partial pressure relation to mol fraction
fraction of total pressure each gas contributes
185
Calculate amount of an insoluble gas collected over wate
Vapor pressure: partial pressure of water, constant at a particular temperature
186
What is the charge on these transition metals?
Zn: zinc
Ag: silver
Cd: cadmium
Zn: 2+ zinc
Ag: +1 silver
Cd: +2 cadmium
187
Which elements are diamtomic? Never found by themselves in nature
188
Identify alkali metals, halogens, noble gases
189
Are all gases diatomic?
Monoatomic gases:
- Noble gases: He, Ne, Ar, Kr, Xe, Rn, Uuo
- Mercury (Hg)
All other gases: diatomic
190
190
First law thermodynamics
Change in internal energy equation: ∆E = ?
∆E = Q + W
Heat (Q), work (W)
191
Sign conventions from system's point of view:
Q
W
+Q: system absorbs heat
-Q: system releases heat
+W: system had work done it
-W: system did work
192
Formation reaction
Rxn that forms 1 mole of substance from its constituent elements at standard state
193
First & second ionization energies equation
194
First & second electron affinity equation
195
electron pooling
in METALS
metallic bonding, atoms share a "pool" of electrons that are free to move throughout the structure, giving rise to properties like electrical conductivity and malleability.
196
Radius trends in isoelectronic species (same number electrons)
as the positive charge increases, the ionic radius decreases
electron drawn closer to the nucleus due to a stronger electrostatic attraction
197
Lattice energy vs bond energy
Lattice energy: energy to pull ions apart [kJ/mol]
Bond energy: breaking up covalent bond
198
Which geometries are planar?
Which geometries are non-polar (when symmetric)?
Planar: linear, trigonal planar, square planar
Non-polar:
- 0 L.P: linear, trigonal planar, tetrahedral. trigonal bipyramidal, octahedral
- 3 more (△): square planar, linear, linear
199
Which atom goes in center of lewis structure?
Atom likes to form the most bond (fewest number of v.e)
Never H
200
Relationship between density (ρ) and molar mass (M) of an ideal gas?
ρ = PM / RT
201
Kinetic molecular theory: ideal gas
Average KE of gas is directly proportional to kelvin temperature
Sample gas, many particles moving in straight line paths in random direction
Pressure gas due to collision gas particles walls container
Real volume of gas particles can be assumed zero (negligible)
- Distance between particles >> size of particles
Gas particles don’t attract/repel each other (no net loss in KE when particles collide)
202
At a given temperature, all gas molecules have the same ?
At a given temperature, all gas molecules have the same average kinetic energy.
203
What is the avg KE for 1 mole of gas?
204
Maxwell Boltzman Velocity Distribution
205
root-mean-square speed (RMS)
206
Graham’s law of effusion
Diffusion: two gasses mix randomly (bidirectional)
Effusion: gas escapes through a pinhole into a vacuum (unidirectional)
207
Exothermic
Endothermic
EXOTHERMIC (-∆H): energy flows from system → surroundings
- Cause surroundings feel hot
ENDOTHERMIC (+∆H): energy flows from surroundings → system
- Cause surroundings feel cold
208
Closed system:
Open system:
Isolated system:
Closed system: matter can’t flow but energy can flow → mass remains constant over time (walls made of conducting material)
Open system: both matter and energy can flow
Isolated system: neither matter or energy can flow
209
Internal energy (∆U)
∆U = KE + PE
210
Heat (Q) vs work (W)
Heat (Q): transfer energy via temp difference (hot → cold)
Work (W): energy transfer when object moved by force (work done on system IE increases)
211
First law thermodynamics
Change in internal energy of system: ∆U = Q + W
+∆U: Heat absorbed by system, work done on system
- ∆U: heat released by system, work done by system
212
Pressure volume work: gas inside rigid cylinder with movable system
1 Latm = 101.3 J
213
Change in enthalpy (∆H):
change in energy for the rxn
energy stored in all bonds of products - all bonds reactants
214
Enthalpy:
At constant P, heat released/absorbed = ?
At constant P, heat released/absorbed = change in enthalpy
q = ∆H
215
Heat capacity (C):
What is main equation?
heat energy required increase the temp of a substance by 1ºC
q = mc∆T
216
Specific heat capacity
amount heat required to raise temp of 1g by 1ºC
specific heat = heat capacity / mass
Intensive property: independent of amount
217
Molar heat capacity (Cn)
amount heat required to raise temp of 1 mole substance by 1ºC
Intensive property: independent of amount
218
Determine enthalpy change of an aqueous reaction:
Constant pressure calorimeter device: measure heat absorbed/released in rxn
219
calorimeter device explination
Heat released by rxn = enthalpy change = heat absorbed by water
Qsolution = mc∆T
(at constant pressure)
220
Determine specific heat of unknown substance: CALORIMTRY
221
Thermochemical equation
chemical equation includes value of ∆H
- Find limiting reactant
- Molar ratio fractions also apply to H
222
Other ways to calculate enthalpy change of rxn:
other than using tabulated values
Hess’s law
Standard enthalpy of formation
Bond enthalpy
223
Standard enthalpy of formation
change in enthalpy when 1 mole compound is formed from its pure elements with all substances in their standard states [kJ/mol]
224
Hess's law
if rxn takes place in several steps, enthalpy change overall reaction = sum of enthalpy changes of individual steps
rxn conditions must be same for each step
225
Use bond enthalpies to estimate enthalpy change rxn
Bond enthalpy: change in enthalpy when 1 mol of covalent bonds in a gaseous compound are broken down to form gaseous products
Higher bond energy = greater energy required to break bond = greater energy released when bond formed
Bond enthalpies are averages (approximate)
Can only use bond enthalpies for rxn where all reacts & products are GASSES
226
1ml = 1 ?
1ml = 1cm^3
227
What is a "salt"?
Combination of ions
228
Rule for determining strong acid
As lin
229
Predicting weak vs strong acid: HF vs HCl
HF (weak acid):
- F small grabs very tightly onto H
- Weak acids hold on to H+
HCl (strong acid):
- Cl bigger, more likely give off H
- Strong acids give off H+
230
Who has a stronger repulsive force bonded groups or lone pairs?
lone pairs
231
Higher lattice energy =
stronger ionic bond
higher melting point
harder
232
In a double replacement reaction, if all ions are soluble (aq)
no reactions forms
(aq) + (aq) → (aq) + (aq) NO REACTIONS
233
Molality
moles solute / mass solvent [mol/kg]
234
Non ideal behavior of gasses:
high P, low T
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