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unit 1 Flashcards

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1
Q

what is an element

A

pure substance made of 1 type of atom

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2
Q

what is a compound

A

pure substance made of 2+ elements chemically bonded (ratio of elements must always be the same)

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3
Q

homogenous mixture

A

uniform composition + properties

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4
Q

heterogeneous mixture

A

non-uniform composition + properties

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5
Q

what’s fractional distillation

A

heating up mixture (of liquids) until 1 or more “fractions” of mixture vaporize

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6
Q

what’s chromatography

A

solutes (stationary) are distributed by liquid/gas (mobile), different solutes = different distribution

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7
Q

what’s gel electrophoresis

A

electrical current is applied to mixture separating charged and uncharged components

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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
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9
Q

particle theory of matter

A
  1. all matter is made of particles
  2. all particles of one substance are identical
  3. particles of matter are in constant motion
  4. particles have spaces b/w them
  5. there are attractive forces b/w particles
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10
Q

sublimation

A

solid –> gas

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11
Q

deposition

A

gas –> solid

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12
Q

ionization

A

gas –> plasma

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13
Q

de-ionization

A

plasma –> gas

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14
Q

what are endothermic processes? give examples

A
  • process requiring/absorbing energy
  • feels cold to the touch
    ex. evaporation, melting, sublimation
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15
Q

what are exothermic processes? give examples

A
  • process producing/releasing energy
  • feels hot to the touch
    ex. freezing, condensation, deposition
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16
Q

what is kinetic molecular theory(KMT)?

A
  • model used to explain/predict behaviour of gases at microscopic level
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17
Q

what postulate/assumptions is the KMT based on?

A
  1. gases are made of tiny particles separated by large distances, most is empty space
  2. gaseous particles are constantly moving in straight lines in random directions
  3. gaseous particles undergo elastic collisions w/ each other + container. loses no kinetic energy
  4. no force of attraction b/w gaseous particles
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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
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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
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20
Q

what does temp measure

A

the average kinetic energy of particles

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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)
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22
Q

what temperature is absolute zero?

A

-273.15C, 0K, -459F

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23
Q

what is water’s melting/freezing point?

A

0C, 273.15K, 32F

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24
Q

what is water’s boiling point?

A

100C, 373.15K, 212F

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25
Q

extrapolation

A
  • estimation of extension of graph/values, based on existing trends
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26
Q

what is the atomic mass unit (AMU)

A
  • 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
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27
Q

what are isotopes

A
  • atoms with the same atomic # but different atomic masses (same # of protons, different # of neutrons)
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28
Q

all isotopes have ____ chemical properties, but _____ physical properties

A

same, slightly different

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29
Q

radioactive decay

A
  • 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
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30
Q

relative atomic mass

A
  • mass of element on periodic table
  • average mass of all isotopes of an atom present on earth
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31
Q

properties of EM spectrum

A
  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)
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32
Q

how are atoms excited

A

when EM radiation is passed through atoms, some is absorbed and used to excite atoms into higher energy

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33
Q

absorption spectrum

A

when wavelengths of light are absorbed by atoms, showing up as dark bands on ROY G BIV spectrum

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34
Q

emission line spectrum

A

when high voltage is applied to the gas, the emission line spectrum is produced
-distinct lines at specific wavelengths ocrrosponding to different elements/compounds

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35
Q

emission spectrum

A

refers to spectrum of light emitted by a source
- includes both conitnuous and line spectra

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36
Q

ground state (stationary state)

A
  • electrons occupying fixed circular orbits around nucleus
  • do not emit energy
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37
Q

how do electrons go from ground state to excited state?

A
  • emitting or absorbing specific amount of energy thats exactly equal to the difference b/w 2 states
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38
Q

electrons ____ energy when going from excited state to ground state and ____ energy when going from ground state to excited state

A

release, absorb

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39
Q

photons of UV radiation have ____ energy than that of infrared radiation

A

more

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40
Q

relationship between photons and energy

A
  • energy of photon = frequency of radiation
  • energy of photon of light = change in energy in atom
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41
Q

Planck equation

A

equation: change in electron energy = planck constant x frequency (E=hv or E=hf)

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42
Q

reds have the ____ photon energy and violets have the ____ photon energy

A

lowest, highest

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43
Q

reds have the ___ wavelength and violets have the ____ wavelength

A

longest, shortest

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44
Q

hydrogen emission spectra

A

red (656nm), blue-green (486nm), blue-violet (434nm), violet (410nm)

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45
Q

what does energy level 1 produce?

A

ultraviolet light

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46
Q

what does energy level 2 produce?

A

visible light

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47
Q

what does energy level 3+ produce?

A

infrared light

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48
Q

what did prince louis de broglie do?

A
  • 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)
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49
Q

what must happen for an object to have a wavelength

A

it must be moving

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50
Q

young’s double slit experiment

A
  • 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
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51
Q

davisson-germer experiment

A
  • 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
52
Q

the uncertainty principle

A
  • we cannot know both the location and momentum of an electron, measuring one blurs the other
53
Q

schrodinger’s wave mechanical model

A
  • 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
Q

what is an atomic orbital

A

region around an atomic nucleus where this is a 90% probability of finding an electron

55
Q

order of sublevels

A

s<p<d<f

56
Q

quantum numbers

A
  • 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
Q

principle quantum number

A
  • represented by n
  • describes size + energy of orbital
58
Q

secondary quantum number(angular momentum)

A
  • represented by l
  • describes number of sub levels
  • describes shape of each orbital
  • corresponds to s,p,d,f
    l = n -1
59
Q

magnetic quantum number

A
  • represented by m(subscript) l
  • describes number of orbitals and orientation within a subshell
60
Q

spin quantum number (fourth quantum number)

A
  • represented by m(subscript) s
  • shows direction of electron spin
  • arrows used in orbital diagrams
61
Q

aufbau principle

A

electrons occupy lowest energy orbital of lowest energy level first

62
Q

pauli’s exclusion principle

A

no two electrons may have same set of quantum numbers therefore electrons in same orbital must have opposite spins

63
Q

hund’s rule

A

distribute electrons in orbitals of equal energy so that no electron pairing occurs until needed

64
Q

what are degenerate orbitals?

A

orbitals of equal energy (ex. 2px, 2py, 2pz)

65
Q

van der Waals

A

atoms that aren’t joined by a bond
- general term to define intermolecular forces b/w atoms

66
Q

internuclear distance

A

the distance b/w the nuclei of adjacent atoms

67
Q

probability of finding electrons in the atomic radius

A
  • probability of finding an electron decreases with increasing distance from nucleus
  • never reaches probability of 0, no outer boundary to atom
68
Q

formula to find approximate attractive force felt by valence electrons from nucleus

A

Zeff = Z - S
Z = nuclear charge = to # of protons in nucleus (atomic #)
S = # of core electrons (total# of electrons - valence)

69
Q

shielding effect

A
  • 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
Q

how does atomic radii change within a group

A

it increases as number of occupied energy levels increase

71
Q

how does atomic radii change within a period

A
  • number of energy levels stay the same however as nuclear charge increases, attraction b/w nucleus and valence electrons increase, decreasing atomic radii
72
Q

what is the ionic radius

A

how size of an atom changes when electrons are added/removed

73
Q

regular atom size vs. cation size

A
  • 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
Q

regular atom size vs. anion size

A
  • 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
Q

what are first ionization energies

A

measure of attraction between nucleus and outer electrons
- energy required to completely remove electron from neutral atom

76
Q

how does IE change within a group

A
  • decreases down a group
  • electron is removed from further away, meaning there’s less attraction to nucleus, requiring less energy
77
Q

how does IE change within a period

A
  • 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
Q

how would first ionization energies be represented in a graph

A
  • 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
Q

what elements have ground states of n=1

A

hydrogen + helium

80
Q

how are atoms ionized

A
  • 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
Q

equation linking frequency, wavelength and speed of light

A

C = fλ
c = speed of light
f = frequency
λ = wavelength

82
Q

equation linking energy, Planck’s constant and frequency

A

E = hf
E = energy
h = Planck’s constant
f = frequency

83
Q

ionization energies are always ____

A

positive

84
Q

second ionization energy

A
  • when another mole of electrons is removed from 1+ ions, turning it into 2+ ions
85
Q

chemical equation of first ionization energy vs. second ionization energy

A

X(g) –>X^+(g) +e^-
X^+(g) –> X^2+(g) + e^-

86
Q

successive ionization energies

A
  • the continued removing of elections until only the nucleus is left
87
Q

trends of successive IE within element

A
  • 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
Q

electron affinity definition

A
  • energy change that occurs when one mole of electrons is added to one mole of gaseous atoms
  • exothermic process
89
Q

electron affinity formula

A

X (g) + e^- –> X^- (g) + energy

90
Q

electron affinity trend across periods

A
  • 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
Q

electron affinity trend across groups

A
  • 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
Q

electronegativity definition

A

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
Q

electronegativity trends across a period

A
  • values increase across period as the effective nuclear charge and atomic radii decrease
94
Q

electronegativity trends down a group

A
  • values decrease down group as the atomic radii increases
  • effective nuclear charge increases however theres more core electrons, shielding effect
95
Q

core electrons

A

electrons of atom, not including valence electrons

96
Q

metallic character across periodic table

A
  • elements get more metallic closer to francium
  • elements get more non-metallic closer to fluorine
97
Q

chemical reactivity across periodic table

A

metals: cr increases down a group, decreases across period
non-metals: cr decrease down a group, increases across period

98
Q

naming polyatomic ions

A

1 extra oxygen: per__ic acid
normal: __ic acid
1 less oxygen: __ous acid
2 less oxygen: hypo__ous acid

99
Q

what state of matter is first/second ionization measured in? why?

A

gas state, least attraction
- in a solid or liquid state, intermolecular forces would affect values

100
Q

what is an oxide

A

made from combination of an element with oxygen

101
Q

ionic character of oxides when going across period 3

A
  • 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
Q

metal oxides (ionic/covalent) combine w/ water to form_____. which ones are weak bases? strong bases?

A

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
Q

non-metal oxides (ionic/covalent) combine with water to form ____. which are strong? weak?

A

covalent, acids (acidic solutions)
ex. P4O10 (s) + 6H2O (l) –> 4H3PO4 (strong/medium acid)
ex. SO3 (g) + H2O (l) –> H2SO4 (strong acid)

104
Q

how is the oxide of silicon different?

A
  • silicon doesn’t dissolve in water
  • classified as acidic oxide since it can neutralize a base
105
Q

amphoteric

A

substances that can react as both an acid and a base

106
Q

aluminum acting as an acid

A

Al2O3 (s) + base –>basic salt + water
- because aluminum can neutralize a base, it acts like an acid

107
Q

aluminum acting as a base

A

Al2O3 (s) + acid –> acidic salt + water

108
Q

acid/base nature going across period 3 of periodic table

A

elements go from very basic to very acidic from left to right

109
Q

physical and chemical properties of first 3 alkali metals

A

phys: good conductors, low density, grey shiny surfaces before oxidizing with oxygen
chem: very reactive, forms ionic compounds w non-metals

110
Q

properties of lithium

A
  • soft + reactive
  • exposure to oxygen gives a dark oxide coat
111
Q

properties of sodium

A
  • softer than lithium
  • more reactive that lithium
112
Q

properties of potassium

A
  • softer them sodium
  • more reactive than sodium
113
Q

reactivity of group 1

A
  • forms single charge ions
  • low ionization energy, shows ease in which outer electron is lost
  • reactivity increases down group
114
Q

first 3 group 1 elements reacting with water

A
  • 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
Q

physical and chemical properties of halogens

A

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
Q

properties of chlorine, bromine, iodine

A
  • 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
Q

reactivity of group 17

A
  • readiness to accept electrons
  • nuclei have high effective charge
  • fluoride is most reactive non-metal, reactivity decreases down group
118
Q

group 1 + group 17 reactions with each other

A
  • 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
Q
A
120
Q

what are halides

A
  • a hydrogen atom bonded with another element
121
Q

melting points of group 1

A
  • 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
122
Q

melting points of group 17

A
  • 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
123
Q

properties of noble gases

A
  • colourless, monatomic, very unreactive (because of inability to lose/gain electrons)
124
Q

why can’t noble gases become cations

A
  • doesn’t form cations since they have highest ionization energies, difficult for electron to be removed
125
Q

why can’t noble gases become anions

A
  • doesn’t form anions because new electrons would be added to empty outer shell, experience negligible effective nuclear for

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