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Tufts - Chem 1 > Exam 1 > Flashcards

Exam 1 Flashcards

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1
Q

4 states of matter for molecules:

SOLID

A

Solid: fixed position, vibrate, can’t leave position in crystalline lattice (defined shape)

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2
Q

4 states of matter for molecules:

LIQUID

A

Liquid: glued together, can slide past each other, vibrate & rotate (shape of container)

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3
Q

4 states of matter for molecules:

GAS

A

Gas: free to move, occupy entire volume container (shape & volume of container)

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4
Q

4 states of matter for molecules:

PLASMA

A

Plasma: “ionized gas” contains electrically charged particles (lightning strikes, TV)

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5
Q

Allotrope

A

Allotropes: different forms of same element

ie. oxygen gas (O2) –> ozone gas (O3)

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6
Q

Pure substance:

A

Pure substance: constant & uniform composition

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7
Q

Element:

A

Element: molecules contain only 1 type of atom

(pure sub. cannot be broken down into simpler substances by chemical changes
O2

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8
Q

Compounds:

A

Compounds: combining different atoms

(pure sub. can be broken down into simpler substances by chemical changes)

H2O

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9
Q

Mixture:

A

Mixture: 2+ types of matter (molecules) that can be present in varying amounts and can be separated by physical changes (ie. evaporation)

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10
Q

Homogenous mixture:

A

Homogenous mixture: uniform composition, appears visually same throughout

(“solution” ie. air)

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11
Q

Heterogeneous mixture:

A

Heterogeneous mixture: composition varies from point to point

distinct clumps different molecules/substance - oil & water

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12
Q

Table used to classify matter:

A
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13
Q

Physical properties:

A

Physical properties: can be observed without changing a substance into another (reversible)

mass, volume, density, boiling pt, solubility, color, softness, something melts

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14
Q

Intensive physical property

A

Intensive: independent of amount of substance present

boiling pt, density, color

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15
Q

Extensive physical property

A

Extensive: depend on amount substance present

weight, mass, length

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16
Q

Chemical properties:

A

Chemical properties: observed when matter undergoes changes in chemical composition

flammability, corrosiveness, reactivity with acid

(hint look for terms with “reacting, changing, burning”)

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17
Q

Changes in matter: Physical change

A

Physical change: don’t change composition of substance and no new substance is formed (wax melts, magnetizing solids)

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18
Q

Changes in matter: chemical change

A

result in formation of new substance with different chemical properties (combustion, oxidation)

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19
Q

Law of conservation of matter:

A

Law of conservation of matter: there is no detectable change in total quantity of matter present when matter converts from one type to another (chemical change), of changes among solid, liquid, gaseous states (physical change)

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20
Q

separation of mixtures techniques

A

Filtration: liquid separated from a solid

Substances with diff solubility can be separated using suitable solvent (sand & salt)

Substances with diff boiling points separated using distillation or evaporation

Sublimation: direct conversation from solid → gas (ammonium chloride)

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21
Q

Signifiant figure rules

A
  • All non-zero digits = significant
    Zeros:
  • Left = not significant
  • Middle = significant
  • Right = significant after decimal point
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22
Q

Rounding number rules

A

Adding/subtracting: same # decimal places as the # with least decimal places

Multiplying/dividing: same # of significant figures 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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23
Q

What can be used to show a larger number of sig figs?

A

scientific notation

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24
Q

All measurements have some degree of uncertainty, not exact. What are the only EXACT numbers?

A

counting, definition, unit conversion (infinite # sig figs, no uncertainty)

They don’t limit # sig figs in a calculation

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25
Q

Accuracy vs. Precision

A

Accuracy: very close to the true/accepted value

Precision: similar results when repeated in the same manner

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26
Q

Density

A

density = mass / volume

At particular temp & pressure, density of substance is characteristic property → often used to identify unknown substance

Determine density of irregular object: use volume water displaced in beaker

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27
Q

Atom:

A

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

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28
Q

Element:

A

1 type of atom

mass is a characteristic feature that is the same for all atoms of that element

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29
Q

Molecule:

A

Molecule: 2+ atoms joined by chemical bonds (could be same atom or different)

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30
Q

Dalton’s atomic theory:

A

wrong, but laid foundation for future work

  1. All matter is made up of tiny particles called atoms (indivisible & indestructible)
  2. Atoms of an element are identical in size, mass, chemical properties
  3. Atoms combine to form compounds in whole number ratios

4.Atoms of element cannot change into atoms another element (only rearrange)

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31
Q

Which chemical reactions are not possible according to Dalton’s atomic theory?

A

Dalton believed: atoms of a given element retain their identities in chemical reactions

CCl4 –> CH4 (not possible)

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32
Q

Cathode ray tube, J.J. Thomson:

A

Early experiment: showed electrons small negatively charged particles inside atom –> plum pudding model

Later experiment using electric filed & magnet: measure charge to mass ratio of electron

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33
Q

Oil drop experiment, Millikan:

A

measure charge on small droplets of oil by suspending them between pair of electrically charged plates

charge of oil droplet are multiples of the electron charge : e-=1.60210-19C

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34
Q

Alpha-ray scattering, Rutherford:

A

alpha particles and gold foil → discovery of nucleus in atoms (disproved the plum pudding model → let to development of modern atomic model)

Atoms much larger than nuclei & mostly empty space inside atom occupied by electrons

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

Useful in determining the nuclear charge of the atom b/c revealed most of atom’s mass & + charge concentrated in nucleus

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35
Q

In a neutral atom, where does most of the mass come from?

A

mass atom comes from protons & neutrons, mass electrons is negligible

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36
Q

Rutherford’s model of atom:

A
  1. All + charge & mass concentrated inside nucleus (tiny region)
  2. Negatively charged particles revolve around nucleus in circular path
  3. Electrostatic force attraction between proton & electrons holds atom together
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37
Q

What were limitations to the Rutherford model of atom?

A

Failed to explain:
1. Stability of an atom

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38
Q

What did Niels Bohr study?

A

electromagnetic radiation

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39
Q

Atomic number (Z):

A

protons in nucleus (found in periodic table)

Neutral atom: electrons = protons

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40
Q

Mass number (A):

A

A = protons + neutrons

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41
Q

Nuclear symbol:

A

represents nucleus of an isotope

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42
Q

Mole:

A

number of atoms/molecules in a bulk sample of matter

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43
Q

Molar mass:

A

Molar mass: mass in grams of 1 mole of that substance [gmol]

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44
Q

Can the number of protons and neutrons in the nucleus of an atom vary?

A

Number or protons in the nucleus defines the element therefore is the same for all atoms of an element. However, the number of neutrons in an atom can vary → isotopes.

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45
Q

Isotopes:

A

same #protons, different #neutrons

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

carbon 12, 13, 14

Isotopes have same chemical but different physical properties (due to differences in mass)

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46
Q

Percentage abundance of isotopes

A
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47
Q

Atomic mass of isotopes

A

weighted average of isotopic masses of all the naturally occurring isotopes of an element (decimal value)

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48
Q
A
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49
Q

Electromagnetic (EM) radiation:

A

oscillating electric & magnetic field perpendicular to each other & direction propagation (ie. visible light from sun, microwaves, x-rays)

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50
Q

Characteristics of EM radiation: wavelength (λ)

A

distance between 2 consecutive peaks/troughs

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51
Q

Characteristics of EM radiation: frequency (ν)

A

cycles pass through given point / second

1/sec = Hz

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52
Q

Amplitude (A):

A

height of peak, corresponds to brightness/intensity

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53
Q

What speed do all types of EM radiation travel?

A

speed of light

c = 3 x 10^8 m/s

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54
Q

Relationship between frequency & wavelength

A
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55
Q

Electromagnetic spectrum: order decreasing λ

A

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

MIVUXG

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56
Q

Memorize wavelength range of visible light:

A

ROYGBIV

red (largest λ) –> violest (smallest λ)

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57
Q

Photoelectric effect (equation)

A

light wave is particulate in nature, consisting of small packets of energy called photons

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58
Q

Photoelectric effect (experiment & findings)

A

Electrons can be ejected from surface of a metal when light have a frequency greater than some threshold shone on it

light with > threshold frequency, KE of emitted electrons increased linearly with frequency of light

KE of emitted electrons didn’t change as intensity light increased

electrons emitted directly proportional to intensity light

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59
Q

Photoelectric effect: threshold frequency (v0)

A

min frequency of light needed to eject electrons

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60
Q

Relationship between threshold frequency and work function

A
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61
Q

Photoelectric effect: max KE of emitted electrons

A
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62
Q

Graph of work function vs. frequency of light (photoelectric effect)

A
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63
Q

Continuous spectrum:

A

contains all wavelengths of visible light

64
Q

Discontinuous spectrum:

A

missing/discontinuous wavelength (line spectra)

65
Q

Line spectra (emission):

A

Heat sample of atoms, absorb energy and become excited (unstable → give off absorbed energy in EM radiation/light).

Atoms didn’t absorb (ground state).
Each type of atom has unique emission spectrum → identify atoms (spectroscopy)

Each emission line consists of a single wavelength of light (implies that light emitted by a gas consists of discrete energies)

66
Q

Absorption spectrum:

A

Shine light on samples of atoms, atoms absorb light of unique wavelengths & become excited. Unabsorbed light comes out → pass it into a prism → photodetector.

67
Q

Emission & absorption spectrums are photographic negatives of each other

A
68
Q

Bohr’s model of atom:

A

combines classic & quantum physics

Stationary states, quantization of angular momentum, radius n’th orbital, energy levels

69
Q

Bohr’s model of atom: Stationary states

A

electrons move around nucleus in circular path of fixed radius & energy called orbits (electron cannot live between orbits)

70
Q

Bohr’s model of atom: Quantization of angular momentum

A
71
Q

Bohr’s model of atom:

radius of the n’th orbital of what kind of atom?

A

applicable for mono electronic species (only 1 electron)

72
Q

Bohr’s model of atom: energy of an electron in the n’th orbital of hydrogen like atom

A
73
Q

Bohr’s model of atom:

electron jumps from low –> high orbit what happens?

A
74
Q
A
75
Q

In an emission spectrum of hydrogen, electron jumps from ni to nf. Find energy & wavelength of the emitted photon.

A
76
Q

What is Bohr’s energy equation used to calculate?

A

Energy released when electron jump

77
Q

What is Ryberg’s equation used to calculate?

A

Calculate wavelength of emitted photon when electron jumps from ni to nf

78
Q

Bohr’s explanation of spectral lines:

A
79
Q

Lyman series:

A

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

80
Q

Balmer series:

A

electron jumps from higher energy level → n=2

Transition produces lowest λ : n=infinity → n=2 (greatest energy gap)

Transition produces highest λ: n=3 → n=2 (smallest energy gap)

81
Q

As you move to higher n values how does the energy gap change?

A

energy gap for next shells is smaller as you move to higher n

wavelength of 1st line in Balmer series > first line Lyman series

82
Q

What is the total number of spectral lines possible?

A

n (initial) - 1

83
Q
A
84
Q

Limitations of Bohr’s Atomic Theory

A
  1. Works only for monoelectronic atoms
  2. Didn’t provide any reason for why electrons can only revolve in orbits where angular moment intregral multiples…
  3. Didn’t provide accurate description of the electron’s location in the atom
85
Q

Quantization of angular momentum

A

v: velocity of object

86
Q

Experimental evidence for wave nature of matter:

A

Davisson & Germer: electrons (particles) have a diffraction pattern (characteristic of a wave)

87
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)

88
Q

Shrodinger wave equation:

A

describe electron wave in 3D (Ψ: x, y, z)

89
Q

Wavefunction: Ψ

A

no physical significance, but can be used to determine the distribution of the electron’s density with respect to the nucleus in an an atom,

90
Q

Square of wave function

A

probability of finding an electron at a particular point

91
Q

Solutions to shrodinger equation

A

set of possible wave functions (ψ), corresponding to a set of orbitals with unique energy (E)

92
Q

Energy of electron in hydrogen atom important points:

A

Only dependent on principle quantum number (n)

Energy levels are quantized → can only have certain discrete energy values

Exact same equation for energy obtained using Bohr’s model

93
Q

Orbital:

A

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

94
Q

To define a particular orbital/wave function, how many quantum numbers do we need?

A
95
Q

Principle quantum number (n):

A

size & energy of the orbital

Possible values: + integer

As value of n increases, size & energy of that orbital increases
All orbitals with same n value are in same shell

96
Q

Total number of allowed orbitals with a given n value

A

n^2

97
Q

Angular momentum quantum number (l)

A

shape of the orbital

for a given n, l is between 0 and n-1
(n>l)

Orbitals with same l value are in same subshell

98
Q

Magnetic quantum number (ml):

A

orientation of the orbital

for a given subshell defined by l, ml can be: -l…0…l (l > ml)

For a particular subshell with defined l value, there are (2l+1) orbitals in subshell = possible values of ml

99
Q

Electron spin quantum number (ms):

A

spin of the ELECTRON, says nothing about the ORBITAL

2 different orientations +1/2,-1/2 (↑, ↓)

100
Q

Each orbital has a max of how many electrons?

A

2 electrons

101
Q
A
102
Q

Types of representations of orbitals

A
103
Q

Node:

A

region with zero electron probability of finding electron

104
Q

Equation for total # nodes, radial nodes, and angular nodes

A

total # nodes = n - 1

radial nodes = n - 1 - l

angular nodes = l

105
Q

Possible orientations of 2p orbital? d orbital? (ml)

A
106
Q

Degenerate orbitals:

A

same n → same amount of E

107
Q

Different subshells have different energies. Arrange them in terms of increasing energy

A

s < p < d < f

108
Q

Shielding effect

A

Electron electron repulsion decreases net force of attraction between nucleus and electron being removed. Reduces net + charge that an electron experiences from the nucleus.

Inner electrons shield outer electrons much more than electrons in the same shell shield each other

109
Q

Penetration effect:

A

Ability of an orbital to attract an electron (distance to nucleus)

Closer electron is to nucleus, lower energy associated with orbital

110
Q

Order penetration of orbitals vs orbital energy

A

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

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

111
Q

Electron configuration:

A

describes how electrons of an atom are filled into atomic orbitals

112
Q

Aufbau principle:

A

fill electrons in lowest energy orbitals first

Ground state configuration: lowest energy config

113
Q

Pauli’s exclusion principle:

A

in a given atom, 2 electrons cannot have the same set of 4 quantum numbers

Orbital hold max 2 electrons with opposite spins

114
Q

Hund’s rule (degenerate orbitals):

A

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)

3 orbitals in p subshell have equal energy
5 orbitals in d subshell have equal energy

115
Q

Electron config of cation:

A

loose ns or np electrons that were added last in Aufbau

EXCEPTION Transition elements: loose ns electrons before losing (n-1)d electrons

116
Q

Abbreviated electron configuration:

A

use previous noble gas

117
Q

Valence shell:

A

outermost shell (highest n value)

involved in chemical bonding (determine chemical properties)

118
Q

Periodic table periods vs. groups

A

Horizontal rows: periods (n)

Vertical columns: groups → elements same group have same valence electron config (similar chemical properties)

119
Q

Anomalous electron config:

A

Cr,Mo,Cu,Ag: unexpected, exceptions (cost additional energy)

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

120
Q

Electron configuration table

A
121
Q

As we go across a period →

A

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

122
Q

Atomic size/radius:

A

bond distance used to approx atomic radius b/c atoms don’t have a sharp boundary & extremely small

123
Q

Periodic table trends

A
124
Q

Nuclear charge (Z):

A

protons in nucleus (or magnitude of + charge)

Greater Z = greater attraction force between nucleus & valence electron = smaller atomic radius

125
Q

Shielding

A

electron - electron repulsion = larger atomic radius

Electrons in same shell have poor shielding

126
Q

Effective nuclear charge (Zeff):

A

pull exerted on an outer electron by the nucleus, taking into account electron-electron repulsion

127
Q

Does removing an electron from an atom change the nucleuar charge?

A
128
Q

Why are valence electrons the easiest to remove from an atom?

A

have highest energies, shielded more, and are farthest from nucleus

129
Q

What is the determining factor for atomic radius?

Across a period? Group?

A

decreases as you move from left to right across a period (due to increasing nuclear charge)

increases as you move down a group (due to the increasing number of electron shells)

130
Q

Isoelectronic series:

A

series of atoms/ions with same number of electrons therefore same shielding

(ie. O2-,F-,Na+)

131
Q

Isoelectronic series: compare the radius between negative and positive ions

A

Most negative ion has largest radius, most positive ion has smallest radius

132
Q

Explain the atomic radius trends across period & down group:

A
133
Q

Ionization energy (IE)

A

energy required to remove an electron from a gaseous atom/ion

Always + for neutral atom (energy required)

134
Q

Ionization energy (IE): exceptions

A

Move from 2A to 3A & 5A to 6A decreases IE instead of increase

first valence electron is being added to a p subshell, the full s subshell is able to shield the p subshell from the nucleus, making the first electron in a p subshell easy to remove.

first set of paired electrons is formed within a p subshell, there is a large amount of electron-electron repulsion within that orbital, which makes the fourth electron added to a p subshell easy to remove.

135
Q

How does the 1st IE compare to 2nd and 3rd IE?

A

IE1 < IE2 < IE3 …

136
Q

How can you identify the # valence electrons in an atom using IE?

A

seeing where a big peak in IE occurs

removing an electron from inner shell requires much much more energy than valence shell

137
Q

Electron affinity:

A

how much an atom wants to gain an electron

(likelihood of a neutral atom to gain an electron)

negative: energy is released when an electron is added
positive: energy must be added to the system to produce an anion
zero: process is energetically neutral

Halogens: high negative EA (more likely to gain electron)
Noble gasses: large positive EA (add electron requires energy, unlikely to gain electron)

138
Q

Ionization energy (IE) vs electron affinity (EA) of Mg+

A
139
Q

Chemical bonds:

A

formed to decrease energy → increase stability

  • Atoms only use valence electrons in bonding
  • Energy must be added to break chemical bonds, forming chemical bonds releases energy
140
Q

Ionic bond:

A

transfer of electrons (metal & nonmetal)

large electronegativity difference

141
Q

Ionic crystal lattice:

A

orderly 3D arrangements of cations & anions held together by electrostatic force of attraction

142
Q

Covalent bond:

Molecules:
Bond length:
Bond energy (BE):
Bonding pair:
Bond order:

A

SHARING electrons (nonmetal & nonmetal)

Molecules: smallest unit of covalent compound

Bond length: combined energy of both bonding atoms

Bond energy (BE): depth of the well at bond length → energy required to break bond or energy released when bond is formed

Bonding pair: shared electrons

Bond order: # electrons shared between a pair of atoms

143
Q

Potential energy vs internuclear distance (covalent bond)

A
144
Q

Non polar covalent bond:

A

equal sharing of electron pair (similar electronegativity)

Example: diatomic elements (Cl-Cl)

145
Q

Polar covalent bond:

A

unequal sharing of electrons (closer to one of the atoms) partial negative charge (poles) NM-

146
Q

ELECTRONEGATIVITY:

A

ability of an atom in a molecule to attract shared electron pair to itself (unmeasurable quantity, not absolute)

More electronegative atom has a stronger attraction for the bonding electrons (∂-), less electronegative atom (∂+)

Electronegativity values with respect to F

147
Q

Greater electronegativity (∆EN) = greater polarity of bond

A
148
Q

Naming: Type 1 binary ionic compounds

A

compound only contains 2 types of elements

metal nonmetal + ide

For cation, name = name of element
For anion, name = first part of elements name + suffix-ide

149
Q

Types of metal cations:

A
150
Q

Naming: Type 2 binary ionic compounds

A

contains metal that forms multiple types of cations

metal (charge on metal) nonmetal + ide

Chromium (I) nitride: Cr3N

151
Q

Ionic compounds with polyatomic ions:

A

compound with +1 atom

Dissolved in water, stays together as a single entity

152
Q

Naming: Binary covalent compounds

A

2 types of non metal atoms

prefix element 1 prefix element 2+ide

  • First element in formula treated as cation (name using full element name)
  • Second element treated anion (named similar to anion in ionic compound -ide)

Denote # atoms present use prefixes: mono (1), di (2), tri (3), tetra (4), pentra (5), hexa (6), hepta (7), octa (8)

NO (mononitrogen monoxide → nitrogen monoxide)

153
Q

Lattice energy:

A

energy released when oppositely charged ions come together to form solid compound

154
Q

When lattice energy is formed: what ions attract more slowly & release more energy

A

Smaller ions & ions with greater charges

Charges (q1, q2) have greater effect on LE than atomic radii (r)

155
Q

Why do atoms with high ionization energy tend to form coavelent bonds?

A

so that they do not form ions (won’t easily give up electrons) and electrons are available for sharing.

Tufts - Chem 1 (4 decks)

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