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Tufts - Chem 1 > Final Exam > Flashcards

Final Exam Flashcards

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1
Q

Allotropes

A

Allotropes: different forms of same element

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2
Q

Pure substance:

A

constant & uniform composition

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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

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4
Q

Mixture:

A

2+ types molecules can be separated by physical changes (ie. evaporation)

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5
Q

Mixture: Homogenous vs Heterogeneous

A

Homogenous mixture: uniform composition - solutions & air (“solution” ie. air)

Heterogeneous mixture: composition varies from point to point

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6
Q

Physical properties:

A

observed without changing substance into another, reversible

(mass, volume, density, boiling pt, solubility, color)

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7
Q

Physical properties: intensive vs extensive

A

Intensive: independent amount substance (boiling pt, density, color)

Extensive: depend amount substance (weight, mass, length)

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8
Q

Chemical properties:

A

undergoes change chemical composition

(flammability, corrosiveness, reactivity with acid)

Key words: reacting, changing, burning

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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)

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10
Q

Significant figure rules

A

All non-zero digits = significant

Zeros:
Left = not significant
Middle = significant
Right = significant after decimal point

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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)

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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)

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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

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14
Q

Atom

A

Atom: smallest unit of an element that can participate in chemical change (indivisible)

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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

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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)
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17
Q

Cathode ray tube, J.J. Thomson

A

DISCOVERED: electrons negative, charge to mass ratio electron

Plum pudding model atom

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18
Q

Oil drop experiment, Millikan:

A

DISCOVERED charge of an electron

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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

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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)

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21
Q

Atomic number (Z):

“Nuclear charge”

A

protons in nucleus (found in periodic table)

Neutral atom: electrons = protons = Z

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22
Q

Mass number (A):

A

protons + neutrons

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23
Q

Nuclear symbol:

A

represents nucleus of an isotope

Atomic number (Z): protons
Mass number (A): protons + neutrons

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24
Q

Empirical formula:

A

molecular formula expressed in lowest whole number ratio

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25
Q

Mole:

A

number of atoms/molecules in a bulk sample of matter

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26
Q

Avogadro’s number

A

Number of particles per mole

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27
Q

Isotopes:

A

atoms with same atomic number (Z) but different mass number (A)

carbon 12, 13, 14

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28
Q

Percent abundance of isotope

A
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29
Q

Atomics mass:

A

weighted average of isotopic masses (mass spectroscopy)

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30
Q

Electromagnetic (EM) radiation:

A

oscillating electric & magnetic field perpendicular to each other & direction propagation → quantized (units: photons)

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31
Q

Wavelength (λ):

A

distance between 2 consecutive peaks/troughs

[m, nm]

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32
Q

Frequency (v)

A

cycles pass through given point / second

[Hz = 1/sec]

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33
Q

Amplitude

A

height of peak, corresponds to brightness/intensity

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34
Q

All types of EM radiation travel at what speed?

A

Speed of light

3 x 10^8 m/s

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35
Q

Electromagnetic spectrum:

A

microwave → infrared → visible → ultraviolet → x-ray → gamma ray

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36
Q

Wavelength range of visible light:

A

(LOW frequency) ROYGBIV (HIGH frequency)

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37
Q

Photoelectric effect:

A

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

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38
Q

Wave theory (classical physics):

A

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
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39
Q

Photoelectric effect: problems

A
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40
Q

Continuous vs discontinous spectrum

A

Continuous spectrum: contains all wavelengths of visible light

Discontinuous spectrum: missing/discontinuous wavelength (line spectra)

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41
Q

Line spectra vs absorption spectra

A

Line spectra (emission): atoms release energy

Absorption spectrum: atoms absorb energy

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42
Q

Bohr’s model of atom:

A

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)
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43
Q

Bohr’s model of atom: equation for amount of energy abosrbed/emitted when electron jumps between energy levels

A

When electrons jump lower → high orbit (absorbs energy), higher → lower (emit energy)

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44
Q

Bohr’s model of atom:

Radius of n’th orbit of MONO electronic atoms (1 electron):

A
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45
Q

Bohr’s model of atom:

Quantization of angular momentum

A
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46
Q

What does Ryberg’s equation tell?

A

find wavelength of photon resulting from an electron jumping between energy levels

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47
Q

What does Bohr’s energy equation say?

A

When electron jumps between energy levels, energy of photon emitted = energy transition

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48
Q

Spectral line: Lyman series

A

electron jumps from higher energy level → ground state (n = 1)

UV light

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49
Q

Spectral line: BALMER series

A

electron jumps from higher energy level → n=2

Visual light

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50
Q

Spectral line: Paschen, Bracker, Pfund

A

Paschen series –> n = 3
Bracker series: –> n = 4
Pfund series: –> n = 5

INFRARED light

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51
Q

How do you compare the wavelength of the 1st line in balmer series to that of lyman series?

A

Wavelength of the 1st line in Balmer series > first line of Lyman series

As we move higher n the ∆E for next shells decreases

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52
Q

Heisenberg uncertainty principle (HUP)

A

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)

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53
Q

Bohr’s explanation of spectral lines:

A
  • 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
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54
Q

Shrodinger wave equation

A

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)

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55
Q

Energy of electron in hydrogen atom important points:

A
  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
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56
Q

Orbital:

A

3D space around the nucleus where the probability of finding an electron is max.

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57
Q

What do we need to define a particular orbital?

A

3 quantum numbers

n: shell, size & energy
l: subshell, shape of orbital
ml: orientation of orbital

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58
Q

Principle quantum number (n)

A

(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)

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59
Q

Angular momentum quantum number

A

(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

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60
Q

Hierarchy: shells, subshells, orbitals

A

shells → subshells → orbitals

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61
Q

Magnetic quantum number

For a particular subshell how many possible orbitals are in subshell?

A

(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

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62
Q

Electron spin quantum number

A

(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)

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63
Q

Each orbital contains max how many electrons?

A

max 2 electrons

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64
Q
A
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65
Q

Ways to visualize orbitals: 3D space around nucleus where probability of finding electron is max

A
  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
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66
Q

Node:

Total nodes
Radial nodes
Angular nodes

A

region with zero electron probability

total #nodes = n - 1

radial nodes = n - 1 - l (spherical)

angular nodes = l (plane)

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67
Q

Degenerate orbitals:

A

same n → same amount of E

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68
Q

What does energy & orbital depend on in multielectron atom?

A

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<p<d<f)

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69
Q

Do subshells have the same amount of energy?

A

Different energies (s < p < d < f)

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70
Q

Shielding effect

Effective nuclear charge

A

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

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71
Q

Penetration effect

A

Distance electrons inside a particular orbital is to nucleus

Closer electron is to nucleus, lower energy associated with orbital

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72
Q

Order penetration of orbitals

Order of orbital energy

A

Order penetration of orbitals: s>p>d>f

Order of orbital energy: s<p<d<f<

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73
Q

Electron configuration

A

how electrons of an atom are filled into atomic orbitals

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74
Q

Aufbau principle

Pauli’s exclusion principle

Hund’s rule

A

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)

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75
Q

Electron configuration for transition elements

A

loose s electrons before losing d electrons

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76
Q

Anomalous electron config: unexpected, exceptions (cost additional energy)

A

Completely filled or ½ filled subshell: more stable (electron shifts to a higher energy orbital)

Cr, Mo, Cu, Ag

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77
Q

As we go across a period

A

add proton to nucleus & electron to valence shell with each element

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78
Q

Periodic Properties: Atomic size/radius

A

Down group (↓) valence electrons are in larger orbitals

Across period (←): # protons increases while # shielding electrons remains same therefore pulled closer to nucleus

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79
Q

Isoelectronic series:

A

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)

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80
Q

Ionization energy:

A

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)

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81
Q

What is a way of identifying number valence electrons in atom using IE?

A

Removing an electron from inner shell&raquo_space; valence electron (much more energy)

identify #v.e by seeing where BIG PEAK in IE occurs

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82
Q

Electron affinity:

A

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)

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83
Q

Ionization energy (IE) exceptions

A

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

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84
Q

Why do atoms form chemical bonds?

A
  • Decrease energy
  • Increase stability
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85
Q

Ionic crystal lattice:

A

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

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86
Q

Molecules

A

smallest unit of covalent compound

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87
Q

Bond length

A

combined energy of both bonding atoms is min

distance which the lowest potential energy is achieved

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88
Q

Bond energy (BE):

A

energy required to break bond or energy released when bond is formed

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89
Q

Breaking chemical bonds vs forming chemical bonds, energy?

A

Break chemical bonds (energy must be added, ENDOTHERMIC)

Forming chemical bonds (release energy, EXOTHERMIC)

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90
Q

Bond order:

A

electrons shared between a pair of atoms

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91
Q

Ionic bond:

A

NM & M (transfer of electrons)

EN difference = +1.7

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92
Q

Covalent bond

A

NM & NM (sharing electrons)

EN difference = less 1.7

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93
Q

Non polar vs polar covalent bond

A

Non polar covalent bond: between same atoms - equal sharing electrons

Polar covalent bond: unequal sharing electrons - partial negative charge most EN

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94
Q

ELECTRONEGATIVITY:

A

tendency of atom to attract electrons towards itself

Greater electronegativity (∆EN) = greater polarity of bond

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95
Q

Naming: Ionic compounds

Basic:
Includes transition metal:
Include polyatomic ions:

A

Basic: M NM + ide

Includes transition metal: M (charge) NM + ide

Include polyatomic ions: M + polyatomic ion name

96
Q

Naming: Binary covalent compounds ( 2 types of NM)

A

Prefix NM prefix NM + ide

Prefix: mono (1), di (2), tri (3), tetra (4), pentra (5), hexa (6), hepta (7), octa (8)

97
Q

Lattice energy:

A

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
Q

Lewis Dot Structure rules

A
  • 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
Q

lewis exceptions

A
100
Q

Resonance structures

A

same arrangement of atoms but diff placement of electrons

Resonance hybrid: superposition/avg of resonance structures

101
Q

What is wrong about how lewis structures depict electrons?

A

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
Q

Bond order:

A

electron pairs shared between two atoms

103
Q

Equivalent vs non-equivalent resonance structure

A

same distribution formal charges

Non-equivalent resonance structure: different distribution FC (Don’t equally contribute to resonance hybrid)

104
Q

Formal charge (FC):

A

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
Q

Criteria for choosing greatest contribution to resonance hybrid:

A

Smaller FC are preferable

Same nonzero FC on adjacent atoms not preferable

A more negative FC should reside on more electronegative atom

106
Q

Choose a likely identity for X in these structures:

A

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
Q

Drawbacks of lewis structure:

A

doesn’t explain shape/geometry of molecule (effect properties)

108
Q

Valence Shell Electron Pair Repulsion Theory (VSEPR):

A

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
Q

Bond strength:

Bond size:

A

Strength: Triple bond > Double bond > Single bond

Size: Triple bond < Double bond < Single bond

110
Q

Electron pair geometry vs Molecular geometry

A

Electron geometry: electron pairs around central atom (includes lone pairs & bonding pairs)

Molecular geometry: atoms relative to central atom (excludes lone pairs)

111
Q

VESPR chart

A

Nonpolar highlighted in blue (if all atoms are same)

112
Q

Valence bond theory (VB):

A

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
Q

How to determine strength bond in VB theory?

A

greater orbital overlap (closer nuclei is to bonded electrons) = stronger bond

114
Q

Sigma bonds (σ):

A

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
Q

Pi bond (π):

A

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
Q

Hybridization:

A

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
Q

Hybridization:

triple, double, single bond

A
118
Q

Sp hybridized orbitals on Be in BeCl2 molecule

A

1s orbital + 1p orbital

119
Q

Sp2 hybridization in BH3 atom

A

1s orbital + 2p orbital

120
Q

Sp^3 hybridization in CH4

A

1s orbital + 3p orbital

121
Q

Does resonance influence hybridization?

A

No! arrangement of π bonds involves only the unhybridized orbitals

Hybridization involves only σ bonds, lone pairs of electrons, and single unpaired electrons

122
Q

Polar molecule:

A

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
Q

Nonpolar molecule

A

symmetric distribution, EN = 0, no lone pairs

124
Q

Intermolecular forces:

A

attraction between molecules with partial charges, or between ions/molecules

125
Q

Dipole dipole moment

A

between POLAR molecules

126
Q

London dispersion forces

A

between ALL atoms due to random motion of electrons

-Stronger LDF in larger, heavier molecules
-Branched molecules have weaker LDF than straight chain

127
Q

Hydrogen bond

A

strong dipole–dipole

H & N, O, F (covalent bond)

128
Q

Molecular orbital theory (MO):

A

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
Q

Valence bond theory vs. molecular orbital theory

A
130
Q

Filling molecular orbitals:

A

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
Q

Bonding orbital

A

electron density directly between nuclei

Placing electron in bonding orbital stabilizes molecule b/c between 2 nuclei

132
Q

Antibonding orbital (*)

A

placing electron in non-bonding orbital destabilises

133
Q

Degenerate orbitals:

A

Electron orbitals having the same energy levels

orbitals in 2p subshell: 2px, 2py, 2pz

134
Q

Delocalized π bond

A

π bond extend over 2+ atoms

In lewis structure, this occurs when with resonance structures involving double & triple bonds

135
Q

Molecular orbital diagram:

A

shows energy and number of electrons present in each MO, atomic orbitals from which each MO is formed

136
Q

P atomic orbiatls –> MO orbitals formed

A
137
Q

Bond order

A

If bond order > 0: stable

Higher bond order = stronger bond

138
Q

Paramagnetic:

Diamagnetic:

A

Paramagnetic: has unpaired electrons
(MO theory can explain paramgnetic, lewis can’t)

Diamagnetic: all electrons paired

139
Q

What is avogadro’s number?

A

number atoms or molecules / moled

140
Q

Percent composition:

A

% by mass of each element in compound (defines identity)

141
Q

Chemical formula

Empirical formula

A

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
Q

Empirical formula mass

A

avg atomic masses of all atoms in empirical formula

143
Q

Limiting reactant:

A

completely consumed first, limits amount of product formed

ICE reaction table: initial, change, end

144
Q

Theoretical, actual, percent yield

A
145
Q

Solutions:

Solute:

Solvent:

A

Solutions: homogeneous mixture of 2+ substances

Solute: smaller amount

Solvent: larger amount (ie. water is universal solvent)

146
Q

Ionic compounds in water:

A

Forces of attraction between solute & solvent particles dissolve

Not all ionic substances are soluble in water

Solubility depends on attraction between ions & water

147
Q

Strong electrolytes

A

dissociate completely into ions (conduct electricity)

  • Soluble ionic compounds (NH4Cl)
  • Strong acids & bases (HCl)
148
Q

Weak electrolytes:

A

dissociate into fewer number of ions (don’t conduct well)

  • Weak acids & bases (HF)
  • Insoluble ionic compounds (AgCl)
149
Q

Non electrolytes

A

Non electrolytes: dissolve in water (polar) but don’t produce ions (don’t conduct)

  • Covalent compounds (C12H22O11)
150
Q

Molarity (M)

A

moles of solute in 1L of solution

151
Q

How to prepare 250mL of a 0.200M aqueous solution of sucrose (aq):

A
152
Q

Dillute vs concentrated

A

Dilute: small amount of solute dissolved

Concentrated: large amount of solute dissolved

153
Q

Dillution

A

Dilution: adding additional solvent to a solution

moles doesn’t change during dilution

154
Q

Chemical reaction types:

A

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
Q

Molecular equations

Complete ionic equations

Net ionic equation

A

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
Q

Alkali metal cations:

A

Li, Na, K, Rb, Cs, Fr

157
Q

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.

A
158
Q

Arrhenius definition of acid & base limited to aqueous solutions

A

Arrhenius acid: acid is a substance that produces H+ ions (protons) in water

Arrhenius base: substance that produces OH- ions in water

159
Q

Bronsted-lowry definition acid & base

A

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
Q

Strong electrolytes:

A

dissociate completely in aq solution

Acids & bases can act as electrolytes (conduct electricity)

161
Q

Electrolyte examples

strong acids
weak acids

strong bases
weak bases

A

Strong acid: HCl HBr Hi & as long as you have 2+ O than H

Strong base: B pt table

162
Q

Neutralization reaction:

A

acid reacts with a base to form salt & water

acid (aq) + base (aq)→ salt (aq) + H2O (l)

Net ionic equation: (H+) + (OH-) → H2O

163
Q

Titration:

titrant, analyte

A

analytical technique to determine concentration of unknown acid/base solution

Titrant: solution known concentration

Analyte: solution unknown concentration

164
Q

Equivalence point:

A

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
Q

How is titration done?

A
166
Q

How to determine equivalence point?

A
167
Q

Redox reactions (oxidation reduction reaction):

A

transfer of electron(s) from one species to another

change in oxidation number (ON)

168
Q
A
169
Q

Oxidation number (ON):

A

Any lone element = 0

Assign most EN element first, ON correspond to group #

Hydrogen: +1 (NM), -1 (M)

Sum ON = charge

170
Q

Important characteristics of redox reaction:

A

oxidation & reduction must occur together

171
Q

Oxidation vs reduction

A

OIL RIG

Element that is OXIDIZED undergoes increase in ON (loss electrons)

Element that is REDUCED undergoes decrease in ON (gain electrons)

172
Q

Cation displacement reaction

“Single replacement reaction”

A

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
Q

Atmospheric pressure:

A

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
Q

Gas pressure

A

Manometer device to measure pressure of gas inside container

175
Q

Boyles law

A

P inversely proportional V

P1V1 = P2V2

176
Q

Charles law

A

V proprtional to T (K)

177
Q

Avogadro’s law

A

V proptional to n

equal volumes of any ideal gas contain equal number of particles (moles)

178
Q

Ideal gas law

A

PV = nRT

Temperature in K

179
Q

Standard temp & pressure (STP):

A

0ºC, 1atm

180
Q

Standard molar volume:

A

1 mole of any ideal gas at standard temp & pressure occupies 22.4L

181
Q

Determine molar mass of gas using ideal gas law:

A
182
Q

Dalton’s law of partial pressure

A

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
Q

Mole fraction

A

Value mole fraction between 0 and 1, sum mol fraction of all components must add up to 1

184
Q

Partial pressure relation to mol fraction

A

fraction of total pressure each gas contributes

185
Q

Calculate amount of an insoluble gas collected over wate

A

Vapor pressure: partial pressure of water, constant at a particular temperature

186
Q

What is the charge on these transition metals?

Zn: zinc
Ag: silver
Cd: cadmium

A

Zn: 2+ zinc

Ag: +1 silver

Cd: +2 cadmium

187
Q

Which elements are diamtomic? Never found by themselves in nature

A
188
Q

Identify alkali metals, halogens, noble gases

A
189
Q

Are all gases diatomic?

A

Monoatomic gases:
- Noble gases: He, Ne, Ar, Kr, Xe, Rn, Uuo
- Mercury (Hg)

All other gases: diatomic

190
Q
A
190
Q

First law thermodynamics

Change in internal energy equation: ∆E = ?

A

∆E = Q + W

Heat (Q), work (W)

191
Q

Sign conventions from system’s point of view:

Q

W

A

+Q: system absorbs heat
-Q: system releases heat

+W: system had work done it
-W: system did work

192
Q

Formation reaction

A

Rxn that forms 1 mole of substance from its constituent elements at standard state

193
Q

First & second ionization energies equation

A
194
Q

First & second electron affinity equation

A
195
Q

electron pooling

A

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
Q

Radius trends in isoelectronic species (same number electrons)

A

as the positive charge increases, the ionic radius decreases

electron drawn closer to the nucleus due to a stronger electrostatic attraction

197
Q

Lattice energy vs bond energy

A

Lattice energy: energy to pull ions apart [kJ/mol]

Bond energy: breaking up covalent bond

198
Q

Which geometries are planar?

Which geometries are non-polar (when symmetric)?

A

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
Q

Which atom goes in center of lewis structure?

A

Atom likes to form the most bond (fewest number of v.e)

Never H

200
Q

Relationship between density (ρ) and molar mass (M) of an ideal gas?

A

ρ = PM / RT

201
Q

Kinetic molecular theory: ideal gas

A

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&raquo_space; size of particles

Gas particles don’t attract/repel each other (no net loss in KE when particles collide)

202
Q

At a given temperature, all gas molecules have the same ?

A

At a given temperature, all gas molecules have the same average kinetic energy.

203
Q

What is the avg KE for 1 mole of gas?

A
204
Q

Maxwell Boltzman Velocity Distribution

A
205
Q

root-mean-square speed (RMS)

A
206
Q

Graham’s law of effusion

A

Diffusion: two gasses mix randomly (bidirectional)

Effusion: gas escapes through a pinhole into a vacuum (unidirectional)

207
Q

Exothermic

Endothermic

A

EXOTHERMIC (-∆H): energy flows from system → surroundings
- Cause surroundings feel hot

ENDOTHERMIC (+∆H): energy flows from surroundings → system
- Cause surroundings feel cold

208
Q

Closed system:

Open system:

Isolated system:

A

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
Q

Internal energy (∆U)

A

∆U = KE + PE

210
Q

Heat (Q) vs work (W)

A

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
Q

First law thermodynamics

A

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
Q

Pressure volume work: gas inside rigid cylinder with movable system

A

1 Latm = 101.3 J

213
Q

Change in enthalpy (∆H):

A

change in energy for the rxn

energy stored in all bonds of products - all bonds reactants

214
Q

Enthalpy:

At constant P, heat released/absorbed = ?

A

At constant P, heat released/absorbed = change in enthalpy

q = ∆H

215
Q

Heat capacity (C):

What is main equation?

A

heat energy required increase the temp of a substance by 1ºC

q = mc∆T

216
Q

Specific heat capacity

A

amount heat required to raise temp of 1g by 1ºC

specific heat = heat capacity / mass

Intensive property: independent of amount

217
Q

Molar heat capacity (Cn)

A

amount heat required to raise temp of 1 mole substance by 1ºC

Intensive property: independent of amount

218
Q

Determine enthalpy change of an aqueous reaction:

A

Constant pressure calorimeter device: measure heat absorbed/released in rxn

219
Q

calorimeter device explination

A

Heat released by rxn = enthalpy change = heat absorbed by water

Qsolution = mc∆T

(at constant pressure)

220
Q

Determine specific heat of unknown substance: CALORIMTRY

A
221
Q

Thermochemical equation

A

chemical equation includes value of ∆H

  • Find limiting reactant
  • Molar ratio fractions also apply to H
222
Q

Other ways to calculate enthalpy change of rxn:

other than using tabulated values

A

Hess’s law

Standard enthalpy of formation

Bond enthalpy

223
Q

Standard enthalpy of formation

A

change in enthalpy when 1 mole compound is formed from its pure elements with all substances in their standard states [kJ/mol]

224
Q

Hess’s law

A

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
Q

Use bond enthalpies to estimate enthalpy change rxn

A

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
Q

1ml = 1 ?

A

1ml = 1cm^3

227
Q

What is a “salt”?

A

Combination of ions

228
Q

Rule for determining strong acid

A

As lin

229
Q

Predicting weak vs strong acid: HF vs HCl

A

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
Q

Who has a stronger repulsive force bonded groups or lone pairs?

A

lone pairs

231
Q

Higher lattice energy =

A

stronger ionic bond

higher melting point

harder

232
Q

In a double replacement reaction, if all ions are soluble (aq)

A

no reactions forms

(aq) + (aq) → (aq) + (aq) NO REACTIONS

233
Q

Molality

A

moles solute / mass solvent [mol/kg]

234
Q

Non ideal behavior of gasses:

A

high P, low T

Tufts - Chem 1 (4 decks)

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