Exam 1 Flashcards
(155 cards)
1
Q
4 states of matter for molecules:
SOLID
A
Solid: fixed position, vibrate, can’t leave position in crystalline lattice (defined shape)
2
Q
4 states of matter for molecules:
LIQUID
A
Liquid: glued together, can slide past each other, vibrate & rotate (shape of container)
3
Q
4 states of matter for molecules:
GAS
A
Gas: free to move, occupy entire volume container (shape & volume of container)
4
Q
4 states of matter for molecules:
PLASMA
A
Plasma: “ionized gas” contains electrically charged particles (lightning strikes, TV)
5
Q
Allotrope
A
Allotropes: different forms of same element
ie. oxygen gas (O2) –> ozone gas (O3)
6
Q
Pure substance:
A
Pure substance: constant & uniform composition
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
8
Q
Compounds:
A
Compounds: combining different atoms
(pure sub. can be broken down into simpler substances by chemical changes)
H2O
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)
10
Q
Homogenous mixture:
A
Homogenous mixture: uniform composition, appears visually same throughout
(“solution” ie. air)
11
Q
Heterogeneous mixture:
A
Heterogeneous mixture: composition varies from point to point
distinct clumps different molecules/substance - oil & water
12
Q
Table used to classify matter:
A
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
14
Q
Intensive physical property
A
Intensive: independent of amount of substance present
boiling pt, density, color
15
Q
Extensive physical property
A
Extensive: depend on amount substance present
weight, mass, length
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”)
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)
18
Q
Changes in matter: chemical change
A
result in formation of new substance with different chemical properties (combustion, oxidation)
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)
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)
21
Q
Signifiant figure rules
A
- All non-zero digits = significant
Zeros: - Left = not significant
- Middle = significant
- Right = significant after decimal point
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)
23
Q
What can be used to show a larger number of sig figs?
A
scientific notation
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
25
Accuracy vs. Precision
Accuracy: very close to the true/accepted value
Precision: similar results when repeated in the same manner
26
Density
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
27
Atom:
Atom: smallest unit of an element that can participate in chemical change (indivisible)
28
Element:
**1 type of atom**
**mass** is a characteristic feature that is the same for all atoms of that element
29
Molecule:
Molecule: 2+ atoms joined by chemical bonds (could be same atom or different)
30
Dalton’s atomic theory:
**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)
31
Which chemical reactions are not possible according to Dalton’s atomic theory?
Dalton believed: atoms of a given element retain their identities in chemical reactions
CCl4 --> CH4 (not possible)
32
Cathode ray tube, J.J. Thomson:
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**
33
Oil drop experiment, Millikan:
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
34
Alpha-ray scattering, Rutherford:
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**
35
In a neutral atom, where does most of the mass come from?
mass atom comes from protons & neutrons, mass electrons is negligible
36
Rutherford’s model of atom:
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
37
What were limitations to the Rutherford model of atom?
Failed to explain:
1. Stability of an atom
38
What did Niels Bohr study?
electromagnetic radiation
39
Atomic number (Z):
protons in nucleus (found in periodic table)
Neutral atom: electrons = protons
40
Mass number (A):
A = protons + neutrons
41
Nuclear symbol:
represents nucleus of an isotope
42
Mole:
number of atoms/molecules in a bulk sample of matter
43
Molar mass:
Molar mass: mass in grams of 1 mole of that substance [gmol]
44
Can the number of protons and neutrons in the nucleus of an atom vary?
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.
45
Isotopes:
**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)
46
Percentage abundance of isotopes
47
Atomic mass of isotopes
weighted average of isotopic masses of all the naturally occurring isotopes of an element (decimal value)
48
49
Electromagnetic (EM) radiation:
oscillating electric & magnetic field perpendicular to each other & direction propagation (ie. visible light from sun, microwaves, x-rays)
50
Characteristics of EM radiation: wavelength (λ)
distance between 2 consecutive peaks/troughs
51
Characteristics of EM radiation: frequency (ν)
cycles pass through given point / second
1/sec = Hz
52
Amplitude (A):
height of peak, corresponds to brightness/intensity
53
What speed do all types of EM radiation travel?
speed of light
c = 3 x 10^8 m/s
54
Relationship between frequency & wavelength
55
Electromagnetic spectrum: order decreasing λ
microwave > infrared > visible > ultraviolet > x-ray > gamma ray
MIVUXG
56
Memorize wavelength range of visible light:
ROYGBIV
red (largest λ) --> violest (smallest λ)
57
Photoelectric effect (equation)
light wave is particulate in nature, consisting of small packets of energy called photons
58
Photoelectric effect (experiment & findings)
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
59
Photoelectric effect: threshold frequency (v0)
min frequency of light needed to eject electrons
60
Relationship between threshold frequency and work function
61
Photoelectric effect: max KE of emitted electrons
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Graph of work function vs. frequency of light (photoelectric effect)
63
Continuous spectrum:
contains all wavelengths of visible light
64
Discontinuous spectrum:
missing/discontinuous wavelength (line spectra)
65
Line spectra (emission):
**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
Absorption spectrum:
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
Emission & absorption spectrums are photographic negatives of each other
68
Bohr’s model of atom:
combines classic & quantum physics
Stationary states, quantization of angular momentum, radius n'th orbital, energy levels
69
Bohr’s model of atom: Stationary states
electrons move around nucleus in circular path of fixed radius & energy called **orbits** (electron cannot live between orbits)
70
Bohr’s model of atom: Quantization of angular momentum
71
Bohr’s model of atom:
radius of the n'th orbital of what kind of atom?
applicable for mono electronic species (only 1 electron)
72
Bohr’s model of atom: energy of an electron in the n'th orbital of hydrogen like atom
73
Bohr’s model of atom:
electron jumps from low --> high orbit what happens?
74
75
In an emission spectrum of hydrogen, electron jumps from ni to nf. Find energy & wavelength of the emitted photon.
76
What is Bohr's energy equation used to calculate?
Energy released when electron jump
77
What is Ryberg's equation used to calculate?
Calculate wavelength of emitted photon when electron jumps from ni to nf
78
Bohr’s explanation of spectral lines:
79
Lyman series:
electron jumps from higher energy level → ground state (n = 1)
80
Balmer series:
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
As you move to higher n values how does the energy gap change?
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
What is the total number of spectral lines possible?
n (initial) - 1
83
84
Limitations of Bohr’s Atomic Theory
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
Quantization of angular momentum
v: velocity of object
86
Experimental evidence for wave nature of matter:
Davisson & Germer: electrons (particles) have a **diffraction pattern** (characteristic of a wave)
87
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)
88
Shrodinger wave equation:
describe electron wave in 3D (Ψ: x, y, z)
89
Wavefunction: Ψ
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
Square of wave function
probability of finding an electron at a particular point
91
Solutions to shrodinger equation
set of possible wave functions (ψ), corresponding to a set of orbitals with unique energy (E)
92
Energy of electron in hydrogen atom important points:
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
Orbital:
3D space around the nucleus where the probability of finding an electron is max
94
To define a particular orbital/wave function, how many quantum numbers do we need?
95
Principle quantum number (n):
**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
Total number of allowed orbitals with a given n value
n^2
97
Angular momentum quantum number (l)
**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
Magnetic quantum number (ml):
**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
Electron spin quantum number (ms):
**spin of the ELECTRON**, says nothing about the ORBITAL
2 different orientations **+1/2,-1/2** (↑, ↓)
100
Each orbital has a max of how many electrons?
2 electrons
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102
Types of representations of orbitals
103
Node:
region with zero electron probability of finding electron
104
Equation for total # nodes, radial nodes, and angular nodes
total # nodes = n - 1
radial nodes = n - 1 - l
angular nodes = l
105
Possible orientations of 2p orbital? d orbital? (ml)
106
Degenerate orbitals:
same n → same amount of E
107
Different subshells have different energies. Arrange them in terms of increasing energy
s < p < d < f
108
Shielding effect
**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
Penetration effect:
Ability of an orbital to attract an electron (distance to nucleus)
Closer electron is to nucleus, lower energy associated with orbital
110
Order penetration of orbitals vs orbital energy
Order penetration of orbitals: s>p>d>f
Order of orbital energy: s
111
Electron configuration:
describes how electrons of an atom are filled into atomic orbitals
112
Aufbau principle:
fill electrons in lowest energy orbitals first
Ground state configuration: lowest energy config
113
Pauli’s exclusion principle:
in a given atom, 2 electrons cannot have the same set of 4 quantum numbers
**Orbital hold max 2 electrons with opposite spins**
114
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 doublin**g up electrons in that orbital set (down)
3 orbitals in p subshell have equal energy
5 orbitals in d subshell have equal energy
115
Electron config of cation:
loose ns or np electrons that were **added last in Aufbau**
EXCEPTION **Transition** elements: loose ns electrons before losing (n-1)d electrons
116
Abbreviated electron configuration:
use previous noble gas
117
Valence shell:
outermost shell (highest n value)
involved in chemical bonding (determine chemical properties)
118
Periodic table periods vs. groups
Horizontal rows: periods (n)
Vertical columns: groups → elements **same group have same valence electron config (similar chemical properties**)
119
Anomalous electron config:
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
Electron configuration table
121
As we go across a period →
add proton to nucleus & electron to valence shell with each element
122
Atomic size/radius:
bond distance used to approx atomic radius b/c atoms don’t have a sharp boundary & extremely small
123
Periodic table trends
124
Nuclear charge (Z):
protons in nucleus (or magnitude of + charge)
Greater Z = greater attraction force between nucleus & valence electron = smaller atomic radius
125
Shielding
electron - electron repulsion = larger atomic radius
Electrons in same shell have poor shielding
126
Effective nuclear charge (Zeff):
pull exerted on an outer electron by the nucleus, taking into account electron-electron repulsion
127
Does removing an electron from an atom change the nucleuar charge?
128
Why are valence electrons the easiest to remove from an atom?
have highest energies, shielded more, and are farthest from nucleus
129
What is the determining factor for atomic radius?
Across a period? Group?
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
Isoelectronic series:
series of atoms/ions with same number of electrons therefore same shielding
(ie. O2-,F-,Na+)
131
Isoelectronic series: compare the radius between negative and positive ions
Most negative ion has largest radius, most positive ion has smallest radius
132
Explain the atomic radius trends across period & down group:
133
Ionization energy (IE)
energy required to **remove an electron** from a gaseous atom/ion
Always + for neutral atom (energy required)
134
Ionization energy (IE): exceptions
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
How does the 1st IE compare to 2nd and 3rd IE?
IE1 < IE2 < IE3 ...
136
How can you identify the # valence electrons in an atom using IE?
seeing where a big peak in IE occurs
removing an electron from inner shell requires much much more energy than valence shell
137
Electron affinity:
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
Ionization energy (IE) vs electron affinity (EA) of Mg+
139
Chemical bonds:
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
Ionic bond:
transfer of electrons (metal & nonmetal)
large electronegativity difference
141
Ionic crystal lattice:
orderly 3D arrangements of cations & anions held together by electrostatic force of attraction
142
Covalent bond:
Molecules:
Bond length:
Bond energy (BE):
Bonding pair:
Bond order:
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
Potential energy vs internuclear distance (covalent bond)
144
Non polar covalent bond:
**equal sharing** of electron pair (similar electronegativity)
Example: diatomic elements (Cl-Cl)
145
Polar covalent bond:
**unequal sharing of electrons** (closer to one of the atoms) partial negative charge (poles) NM-
146
ELECTRONEGATIVITY:
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
Greater electronegativity (∆EN) = greater polarity of bond
148
Naming: Type 1 binary ionic compounds
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
Types of metal cations:
150
Naming: Type 2 binary ionic compounds
contains **metal that forms multiple types of cations**
metal (charge on metal) nonmetal + ide
Chromium (I) nitride: Cr3N
151
Ionic compounds with polyatomic ions:
compound with +1 atom
Dissolved in water, stays together as a single entity
152
Naming: Binary covalent compounds
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
Lattice energy:
energy released when oppositely charged ions come together to form solid compound
154
When lattice energy is formed: what ions attract more slowly & release more energy
Smaller ions & ions with greater charges
Charges (q1, q2) have greater effect on LE than atomic radii (r)
155
Why do atoms with high ionization energy tend to form coavelent bonds?
so that they do not form ions (won't easily give up electrons) and electrons are available for sharing.
