1
Q
what is an element
A
pure substance made of 1 type of atom
2
Q
what is a compound
A
pure substance made of 2+ elements chemically bonded (ratio of elements must always be the same)
3
Q
homogenous mixture
A
uniform composition + properties
4
Q
heterogeneous mixture
A
non-uniform composition + properties
5
Q
what’s fractional distillation
A
heating up mixture (of liquids) until 1 or more “fractions” of mixture vaporize
6
Q
what’s chromatography
A
solutes (stationary) are distributed by liquid/gas (mobile), different solutes = different distribution
7
Q
what’s gel electrophoresis
A
electrical current is applied to mixture separating charged and uncharged components
8
Q
why does temperature stay the same during phase changes?
A
- kinetic energy remains constant
- energy is used to overcome attractive forces between atoms + disrupt solid lattice
9
Q
particle theory of matter
A
- all matter is made of particles
- all particles of one substance are identical
- particles of matter are in constant motion
- particles have spaces b/w them
- there are attractive forces b/w particles
10
Q
sublimation
A
solid –> gas
11
Q
deposition
A
gas –> solid
12
Q
ionization
A
gas –> plasma
13
Q
de-ionization
A
plasma –> gas
14
Q
what are endothermic processes? give examples
A
- process requiring/absorbing energy
- feels cold to the touch
ex. evaporation, melting, sublimation
15
Q
what are exothermic processes? give examples
A
- process producing/releasing energy
- feels hot to the touch
ex. freezing, condensation, deposition
16
Q
what is kinetic molecular theory(KMT)?
A
- model used to explain/predict behaviour of gases at microscopic level
17
Q
what postulate/assumptions is the KMT based on?
A
- gases are made of tiny particles separated by large distances, most is empty space
- gaseous particles are constantly moving in straight lines in random directions
- gaseous particles undergo elastic collisions w/ each other + container. loses no kinetic energy
- no force of attraction b/w gaseous particles
18
Q
how do gases act at high temp, low pressure?
A
- forces b/w gas molecules are minimized
- high degree of separation
- adheres to ideal gas model
19
Q
how do gases act at low temp, high pressure?
A
- particles move slower
- distance b/w particles decrease
- intermolecular attractions become significant, gas can liquefy
- gas departs from ideal gas behavior, exhibits real gas behavior
20
Q
what does temp measure
A
the average kinetic energy of particles
21
Q
what does the kelvin temperature scale represent
A
- the relationship b/w temperature and volume within gases (experiments showed that changing temp of gas changes its volume)
22
Q
what temperature is absolute zero?
A
-273.15C, 0K, -459F
23
Q
what is water’s melting/freezing point?
A
0C, 273.15K, 32F
24
Q
what is water’s boiling point?
A
100C, 373.15K, 212F
25
extrapolation
- estimation of extension of graph/values, based on existing trends
26
what is the atomic mass unit (AMU)
- relative unit of measure of atomic/molecular weights
- equal to one-twelfth of the mass of an atom of carbon-12(standard)
- around 1.67 x 10^-27
27
what are isotopes
- atoms with the same atomic # but different atomic masses (same # of protons, different # of neutrons)
28
all isotopes have ____ chemical properties, but _____ physical properties
same, slightly different
29
radioactive decay
- some isotopes are stable while others are not (too many/few neutrons) -->go through nuclear decay
- becomes radioactive (radioisotopes)
- leads to spontaneous transformation from 1 isotope/element into another
30
relative atomic mass
- mass of element on periodic table
- average mass of all isotopes of an atom present on earth
31
properties of EM spectrum
1. all EM radiation travels at the same speed in a vacuum (3.00 x 10^8m/s)
2. as wavelength of radiation increases, frequency decreases
3. sunlight and white light produce entire ROY G BIV spectrum
4. amplitude of wave represents intensity (higher amplitude = greater intensity)
32
how are atoms excited
when EM radiation is passed through atoms, some is absorbed and used to excite atoms into higher energy
33
absorption spectrum
when wavelengths of light are absorbed by atoms, showing up as dark bands on ROY G BIV spectrum
34
emission line spectrum
when high voltage is applied to the gas, the emission line spectrum is produced
-distinct lines at specific wavelengths ocrrosponding to different elements/compounds
35
emission spectrum
refers to spectrum of light emitted by a source
- includes both conitnuous and line spectra
36
ground state (stationary state)
- electrons occupying fixed circular orbits around nucleus
- do not emit energy
37
how do electrons go from ground state to excited state?
- absorbing specific amount of energy thats exactly equal to the difference b/w 2 states
38
electrons ____ energy when going from excited state to ground state and ____ energy when going from ground state to excited state
release, absorb
39
photons of UV radiation have ____ energy than that of infrared radiation
more
40
relationship between photons and energy
- energy of photon = frequency of energy
- energy of light emitted from photon= change in energy in atom
41
Planck equation
equation: change in electron energy = planck constant x frequency (E=hv or E=hf)
42
reds have the ____ photon energy and violets have the ____ photon energy
lowest, highest
43
reds have the ___ wavelength and violets have the ____ wavelength
longest, shortest
44
hydrogen emission spectra
red (656nm), blue-green (486nm), blue-violet (434nm), violet (410nm).
45
what does energy level 1 produce?
ultraviolet light
46
what does energy level 2 produce?
visible light
47
what does energy level 3+ produce?
infrared light
48
what did prince louis de broglie do?
- PhD thesis: if things believed to be particles (electrons, cars) could act like waves
- no one took his idea seriously until einstein read them and agreed
- came up with the formula: wavelength = h/mv, allowing the calculation of wavelength of moving particle (m=mass, v=velocity)
49
what must happen for an object to have a wavelength
it must be moving
50
young's double slit experiment
- a light source illuminates barrier with 2 slits
- 2 beams of light can be seem from slits
- as light waves spread, alternating dark and light beams can be seen
- evidence that light behaves as a wave
51
davisson-germer experiment
- shot elections at crystal sample of nickel
- spaces b/w nickel atoms are similar in size to wavelength of moving electrons
- on other side of nickel, electrons (acting as waves) hit screen and formed interference pattern
- conc: electrons act like waves
52
the uncertainty principle
- we cannot know both the location and momentum of an electron, measuring one blurs the other
53
what did schrodinger's wave mechanical model do
- proposed an equation that instead of being able to calculate exact location of electron, gives probability of finding electron in a specific place around nucleus
54
what is an atomic orbital
region around an atomic nucleus where this is a 90% probability of finding an electron
55
order of sublevels
s
56
quantum numbers
- used to describe position of electron
- 4 quantum numbers
- pauli's exclusion principle states no two electrons have the same set of quantum numbers
57
principle quantum number
- represented by n
- describes size + energy of orbital
58
secondary quantum number(angular momentum)
- represented by l
- describes number of sub levels
- describes shape of each orbital
- corresponds to s,p,d,f
l = n -1
59
magnetic quantum number
- represented by m(subscript) l
- describes number of orbitals and orientation within a subshell
60
spin quantum number (fourth quantum number)
- represented by m(subscript) s
- shows direction of electron spin
- arrows used in orbital diagrams
61
aufbau principle
electrons occupy lowest energy orbital of lowest energy level first
62
pauli's exclusion principle
no two electrons may have same set of quantum numbers therefore electrons in same orbital must have opposite spins
63
hund's rule
distribute electrons in orbitals of equal energy so that no electron pairing occurs until needed
64
what are degenerate orbitals?
orbitals of equal energy (ex. 2px, 2py, 2pz)
65
van der Waals
atoms that aren't joined by a bond
- general term to define intermolecular forces b/w atoms
66
internuclear distance
the distance b/w the nuclei of adjacent atoms
67
probability of finding electrons in the atomic radius in relation to distance from nucleus
- probability of finding an electron decreases with increasing distance from nucleus
- never reaches probability of 0, no outer boundary to atom
68
formula to find approximate attractive force felt by valence electrons from nucleus
Zeff = Z - S
Z = nuclear charge = to # of protons in nucleus (atomic #)
S = # of core electrons (total# of electrons - valence)
69
shielding effect
- how electrons will block/shield the nuclear attraction of the nucleus from the valence electrons
- more electrons b/w valence electrons and nucleus = less attractive forces felt by valence electrons
70
how does atomic radii change within a group
it increases as number of occupied energy levels increase
71
how does atomic radii change within a period
- number of energy levels stay the same however as nuclear charge increases, attraction b/w nucleus and valence electrons increase, decreasing atomic radii
72
what is the ionic radius
how size of an atom changes when electrons are added/removed
73
regular atom size vs. cation size
- cations are smaller due to increased nuclear attraction (less electrons that need attraction)
- for isoelectronic cations, the more positive the ionic charge, the smaller the ionic radius
74
regular atom size vs. anion size
- anions are larger due to increase in electron-electron repulsion (same charge = pushes each other away)
- for isoelectronic anions, the more negative the charge, the larger the ionic radius
75
what are first ionization energies
measure of attraction between nucleus and outer electrons
- energy required to completely remove electron from neutral atom
76
how does IE change within a group
- decreases down a group
- electron is removed from further away, meaning there's less attraction to nucleus, requiring less energy
77
how does IE change within a period
- increases across a period
- increase of electron nuclear charge causes increase in attraction b/w outer electrons and nucleus
- electrons are more difficult to remove
78
how would first ionization energies be represented in a graph
- group 1 --> troughs
- group 18--> peaks
- after every energy level shell, there is a drop
- a drop occurs when electrons start pairing (ex. b/w N and O)
- general upwards trend (IE increases across period)
79
what elements have ground states of n=1
hydrogen + helium
80
how are atoms ionized, what is the point of ionization called?
- as electrons move up the levels, they reach the convergence limit (n=∞) and converge, forming continuum
- here, no more energy is needed to promote electron, electron leaves atom, atom is ionized
81
equation linking frequency, wavelength and speed of light
C = fλ
c = speed of light
f = frequency
λ = wavelength
82
equation linking energy, Planck's constant and frequency
E = hf
E = energy
h = Planck's constant
f = frequency
83
ionization energies are always ____
positive
84
second ionization energy
- when another mole of electrons is removed from 1+ ions, turning it into 2+ ions
85
chemical equation of first ionization energy vs. second ionization energy
X(g) -->X^+(g) +e^-
X^+(g) --> X^2+(g) + e^-
86
successive ionization energies
- the continued removing of elections until only the nucleus is left
87
trends of successive IE within element
- successive IE increases for all atoms as remaining electrons experiences more effective nuclear charge because of proximity to nucleus
- jumps when electrons are removed from levels closer to nucleus, as they're exposed to more positive charge from nucleus, needing more energy to be removed
88
electron affinity definition
- energy change that occurs when one mole of electrons is added to one mole of gaseous atoms
- ability of atom to accept an electron
- exothermic process
89
electron affinity formula
X (g) + e^- --> X^- (g) + energy
90
electron affinity trend across periods
- become more negative across period
- halogens have most EA since they only need 1 electron to be isoelectric
- elements in group 15 have less negative EA, half filled orbital
91
electron affinity trend across groups
- down group, in general EA become less negative --> last 3-4 elements have little difference b/w EA values
- group 1 metals have lowest effective nuclear charge, attracts extra electron the least.
92
electronegativity definition
relative attraction that an atom has for the shared pairs of electrons in a covalent bond
- ability of an atom in a molecule to attract electrons to itself
93
electronegativity trends across a period
- values increase across period as the effective nuclear charge increases and atomic radii decrease
94
electronegativity trends down a group
- values decrease down group as the atomic radii increases
- effective nuclear charge increases however theres more core electrons, shielding effect
95
core electrons
electrons of atom, not including valence electrons
96
metallic character across periodic table
- elements get more metallic closer to francium
- elements get more non-metallic closer to fluorine
97
chemical reactivity across periodic table
metals: cr increases down a group, decreases across period
non-metals: cr decrease down a group, increases across period
98
naming polyatomic ions
1 extra oxygen: per__ic acid
normal: __ic acid
1 less oxygen: __ous acid
2 less oxygen: hypo__ous acid
99
what state of matter is first/second ionization measured in? why?
gas state, least attraction
- in a solid or liquid state, intermolecular forces would affect values
100
what is an oxide
made from combination of an element with oxygen
101
acidity/alkinity of oxides across period 3
- oxides get more acidic
- argon doesn't make an oxide (noble gas)
- outliers are SiO2 and Al2O3 which do not react with water, but rather acidic/alkaline solutions
102
metal oxides (ionic/covalent) combine w/ water to form_____. which ones are weak bases? strong bases?
ionic, bases (metal hydroxides)
ex. Na2O (s) + H2O (l) --> 2NaOH (aq), a strong base --> group 1 hydroxides +barium hydroxide
ex. NgO (s) + H2O (l) --> Mg(OH)2 (aq), a weak base
103
non-metal oxides (ionic/covalent) combine with water to form ____. which are strong? weak?
covalent, acids (acidic solutions)
ex. P4O10 (s) + 6H2O (l) --> 4H3PO4 (strong/medium acid)
ex. SO3 (g) + H2O (l) --> H2SO4 (strong acid)
104
how is the oxide of silicon different?
- silicon doesn't dissolve in water
- classified as acidic oxide since it can neutralize a base
105
amphoteric
substances that can react as both an acid and a base
106
equation of aluminum acting as an acid
Al2O3 (s) + base -->basic salt + water
- because aluminum can neutralize a base, it acts like an acid
107
aluminum acting as a base
Al2O3 (s) + acid --> acidic salt + water
108
acid/base nature going across period 3 of periodic table
elements go from very basic to very acidic from left to right
109
physical and chemical properties of first 3 alkali metals
phys: good conductors, low density, grey shiny surfaces before oxidizing with oxygen
chem: very reactive, forms ionic compounds w non-metals
110
properties of lithium
- soft + reactive
- exposure to oxygen gives a dark oxide coat
111
properties of sodium in relation to lithium
- softer than lithium
- more reactive that lithium
112
properties of potassium in relation to sodium
- softer them sodium
- more reactive than sodium
113
reactivity trend of group 1, what makes group 1 reactive?
- forms single charge ions
- relatively low ionization energy, shows ease in which outer electron is lost
- reactivity increases down group
114
first 3 group 1 elements reacting with water
- alkali metal+water --> hydrogen+metal hydroxide
- lithium in water: floats, reacts slowly, releases hydrogen, keeps shape
- sodium in water: vigorous release of hydrogen, heat produced melts unreacted metal forming small ball that moves on water surface
- potassium in water: produces enough heat to ignite hydrogen release by reaction, produces lilac flame
115
physical and chemical properties of halogens
phys: coloured, includes gases(F2, Cl2), liquids(Br2), solids(I2, At2)
chem: very reactive non-metals (decreases down group), forms ionic compounds w/ metals and covalent bonds w/ other non-metals
116
properties of chlorine, bromine, iodine
- all toxic + reactive non-metals
- at room temp, Cl is green gas, Br is dark liquid that turns into dark brown gas, I is a crystalline solid
117
reactivity of group 17
- readiness to accept electrons
- nuclei have high effective charge
- fluoride is most reactive non-metal, reactivity decreases down group
118
group 1 + group 17 reactions with each other
- react to create halides
- halogen atom gains electron from alkali metal
- electrostatic force of attraction bonds ions together
- most reactive reaction occurs with elements furthest apart (ex. francium and fluorine)
119
what are halides
- a halogen atom bonded with another element
120
melting points of group 1
- depends on bond + structure
- group 1: going down the group, melting points decrease
- elements have metallic structure held together by attractive forces, attractive forces decrease as atom gets larger, needing less energy to melt
121
melting points of group 17
- group 17: going down the group, melting points increase
- molecular structures held together by london forces which increases with # of electrons, more electrons in elements down the group
122
properties of noble gases
- colourless, monatomic, very unreactive (because of inability to lose/gain electrons)
123
why can't noble gases become cations
- doesn't form cations since they have highest ionization energies, difficult for electron to be removed
124
why can't noble gases become anions
- doesn't form anions because new electrons would be added to empty outer shell, experience negligible effective nuclear force
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