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Chemistry H > MIDTERM! > Flashcards

MIDTERM! Flashcards

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1
Q

mass vs weight

A
  • mass is a measure of object resistance
  • weight is measured through earth’s gravity
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2
Q

precision vs accuracy

A
  • precision: degree of agreement among several measurements of the asme quantity
  • accuracy: agreement of a particular value with its true value
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3
Q

random error

A

the measured value has an equal chance of being too high or too low (not precise, not accurate)

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4
Q

systematic error

A

the measured value will always high (or low)/ dates off by the same amount each time (precise, not accurate)

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5
Q

sig figs non zero integers

A

always count as significant figures

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6
Q

sig fig zero rules

A
  1. leading zeros - never count (0.0025 is only 2 sig figs)
  2. captive zeros - always count (1.008 is 4 sig figs)
  3. trailing zeros - only count if # has a decimal point
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7
Q

operations with sig figs

A

sig fig in result has the same amount of sig figs as the one with least amount of sig figs

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8
Q

regular notation

A

standard way to weight numbers

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9
Q

scientific notation

A

short handed way of writing (2.8 *10^24315)

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10
Q

giga (G)

A

10^9

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11
Q

mega (M)

A

10^6

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12
Q

kilo (k)

A

10^3

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13
Q

deci (d)

A

10^-1

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14
Q

centi (c)

A

10^-2

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15
Q

milli (m)

A

10&-3

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16
Q

micro (u)

A

10^-6

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17
Q

nano (n)

A

10^-9

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18
Q

pico (p)

A

10^-12

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19
Q

femto (f)

A

10^-15

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20
Q

C = (F-32) * 5/9

A

fahrenheit to celsius

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21
Q

celsius to fahrenheit

A

f = 9/5c + 32

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22
Q

celsius to kelvin

A

k = c + 273

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23
Q

matter classifies into

A
  1. pure substances
  2. mixtures
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24
Q

pure substances classifies into

A
  1. elements (one type of atom)
  2. compound (more than one type of atoms)
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25
Q

mixtures classifies into

A
  1. homogeneous (uniform)
  2. heterogenous (non uniform)
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26
Q

proton basic info

A

+ charge, 1 amu

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27
Q

neutron charge

A

0 charge, 1 amu

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28
Q

electron basic info

A
  • charge, 1/1836 amu
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29
Q

isotopes

A

atoms of the same elements with different #s of neutrons (mass changes)

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30
Q

ions

A

charged atom (gain or lose e-) (mass doesn’t change)

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31
Q

atom box LABELS

A
  1. top right, mass number
  2. top left, charge
  3. bottom left, atomic #
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32
Q

weighted average

A

((relative abundance * mass of isotope) + (relative abundance * mass of isotope 2))/100

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33
Q

proton and electron relation

A

equal, unless there’s a charge

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34
Q

proton and neutron relation

A

mass # - neutrons = protons
nass # - protons = neutrons

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35
Q

aristotle

A

340 BCE
- greek philosopher aristotle thought fire, water, air, and earth were the building blocks of everything

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36
Q

dalton

A

5 part atomic theory
1) matter is made up of atoms that are indivisible and indestructible
2) all atoms of an element are identical
3) atoms of diff elements have diff weights/chemical properties
4) atoms of diff elements combine in simple whole numbers to form compounds
5) atoms cannot created or destroyed

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37
Q

rutherford

A
  • scattering experiment, rutherford sent alpha particles through a thin sheet of gold
  • allowed him to discover nucleus
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38
Q

bohr

A
  • enhanced understanding of atomic structure and quantum theory
  • proposed a model of the atom where electrons were able to occupy only certain orbits around the nucleus
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39
Q

jj thompson

A
  • discovered the electron
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40
Q

chadwick

A
  • discovered atoms not only consist of protons and electrons but also neutrons!
  • neutral subatomic particle has around the same mass as a proton
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41
Q

democritus

A
  • father of modern science
  • atoms are the basic building block of matter

1) all matter consists of atoms, which cannot be further divided
2) atoms are extremely small - too small to see
3) atoms are solid particles that are indestructible
4) atoms are serrated by one another by emptiness or “void”

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42
Q

schrodinger & heisenberg

A
  • schrodinger explored the idea that electrons move more like waves than particles
  • his ideas led Heisenberg to develop the uncertainty principle (if an electron moved as a wave, it would be impossible to simultaneously measure both its position and momentum)
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43
Q

chemical properties

A

describe a substances ability to change to a difference substance

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44
Q

element

A

cannot be broken down chemically into simpler substance

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45
Q

compounds

A

can be broken down chemically into elements (are 2 or more atoms bounded together)

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46
Q

solvent

A

does the dissolving

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47
Q

solute

A

gets dissolved

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48
Q

alloy

A

one or more slides dissolved in another solid

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49
Q

lowest energy level

A

where electron starts from is ground state
- electron configuration written in lowest energy
- atomic spectra

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50
Q

excited state

A
  • heat, electricity, or light can move up to different energy levels
  • when it falls back to ground state, it gives back energy as light
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51
Q

returning to ground state

A
  • may fall down in specific steps
  • each step has different energy
  • the further they fall, the more energy released = higher frequencies
  • orbitals also have different energies inside energy levels
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52
Q

quantum mechanics

A

an explanation of how small particles behave - an explanation for subatomic particles and atoms as waves

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53
Q

classical mechanics

A

describes the motions of bodies much larger than atoms

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54
Q

heisenberg uncertainty principles

A

you cannot know both the position and momentum of an electron (where its going vs where it is)
- warner heisenberg!

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55
Q

ernest rutherford model

A
  • gold foil experiment to discover dense positive piece at nucleus
  • electrons move around like planets around the sun
  • mostly empty space
  • did NOT explain chemical properties of elements
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56
Q

niel bohr’s model

A
  • move like planets around the sun
  • specific circular orbits at different levels
  • an amount of fixed energy separates one level from another
  • electrons can jump from one level to another (circular paths)
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57
Q

ladder rungs

A
  • energy level: measure of fixed energy e-
  • electrons cannot between nergy levels
  • you can’t stand between ladder/rungs”

UNLIKE ladders: rungs are not evenly spread
- higher level are closer together, less energy needed for jump for jump

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58
Q

quantum mechanical medol

A
  • energy is “quantized” into chunks
  • a quantum is is exact energy needed to move e- one energy level to another
  • quantum leaps in energy because e- cannot exist between energy levels
59
Q

principle quantum number (n)

A
  • denotes the energy level (shell) e- is located in
  • maximum number of e- that can fit into the dingy level is 2n^2
    erwin schrödinger (1926) derived an equation that describes energy and position of e- in an atom
60
Q

atomic orbitals

A

for each energy level, rhcrodinger’s equation describes several shapes called atomic orbitals
- only tells probability of finding e- from a certain difference from nucleus, inside blurry cloud

61
Q

sublevels

A

s, p, d, f

62
Q

s (spherical)

A

1 orbital, 2 maximum electrons

63
Q

p (dumbell)

A

3 orbitals, 6 maximum electrons

64
Q

d (clover leaf)

A

5 orbitals, 10 maximum electrons

65
Q

f (complicated)

A

7 orbitals, 14 maximum electrons

66
Q

energy level 1!!

A
  • only s sublevel
  • 1s2 (1 orbital) w/ only 2e-
    2 total e-
67
Q

energy level 2!!

A
  • has s and p sublevels
  • 2s2 (2 orbitals) w 2e-
  • 2p6 (3 orbitals) w 6e-
    2s2 2p6 ~ 8 total e-
68
Q

energy level 3!!

A
  • s, p, d sublevels
  • 3s2 (1 orbital) w 2e-
  • 3p6 (3 orbitals) w 6e-
  • 3d10 (5 orbitals) w 10e-
    3s2 3p6 3d10 ~ 18 total e-
69
Q

energy level 4!!

A
  • s, d, p, f sublevels
  • 4s2 (1 orbital) w 2e-
  • 4p6 (3 orbitals) w 6e-
  • 4d10 (5 orbitals) w 10e-
  • 4f14 (7 orbitals) w 14e-
    4s2 4p6 4d10 4f14 ~ 32 total e-
70
Q

aufbau principle

A

electrons enter the lowest energy first (causes difficulties because overlap of orbitals of different energies)

71
Q

pauli exclusion principle

A

2 electrons max per orbital ~ different spins. no two electrons in an atom have the same 4 quantum numbers

72
Q

hund’s rule

A

when electrons occupy orbitals of equal energy, they don’t air up until they have to

73
Q

chromium exception

A

4s2 3d4 –> 4s1 3d5

74
Q

copper exception

A

4s2 3d9 –> 4s1 3d10

75
Q

group number

A

number of valence electrons of a main group atom = drop number
- atoms like to fill or empty outermost shells (octet rule!)
- outer level contains 2s electrons and 6p electrons

76
Q

atomic size

A
  • increases going down a group
  • decreases going across a period
77
Q

ion size

A
  • cations smaller than the atoms they come from
  • anions larger than the atoms they come from
78
Q

ionization energy

A
  • energy required to remove an electron from an atom
  • IE increases across a period
  • metals lose electrons more easily than nonmetals
  • nonmetals lose electrons with difficulty because they like to gain electrons
79
Q

electronegativity

A
  • measure of the ability of an atom in a molecule to attract electrons to itself
  • electronegativity increases up a group of elements
  • electronegativity increases right in a period of elements
80
Q

metallic character

A

a measure of how easily an atom uses an e-

  • non metallic at F, CL
  • most metallic at Cs, Fr
81
Q

alkali metals

A

tend to form +1 ions

82
Q

alkali earth metals

A

tend to form +2 ions

83
Q

halogens

A

tend to form -1 ions

84
Q

group 1 (+1)

A

li
na
k
pb
cs

85
Q

group 2 (+2)

A

be
mg
ca
sr
ba

86
Q

group 3-12

A

nothing lol

87
Q

group 13 (+3)

A

al

88
Q

group 14

A

nothing lol

89
Q

group 15 (-3)

A

n
p
as

90
Q

group 16 (-2)

A

o
s
se

91
Q

group 17 (-1)

A

f
cl
br

92
Q

group 18 (0)

A

nothing lol

93
Q

h2po4-

A

dihydrogen phosphate

94
Q

c2h3o2

A

acetate

95
Q

hso3-

A

hydrogen sulfite

96
Q

hso4-

A

hydrogen sulfate (bisulfate)

97
Q

hco3-

A

hydrogen carbonate (bicarbonate)

98
Q

no2-

A

nitrite

99
Q

no3-

A

nitrate

100
Q

cn-

A

cyanide

101
Q

oh-

A

hydroxide

102
Q

mno4-

A

permanganate

103
Q

clo-

A

hypochlorite

104
Q

clo2-

A

chlorite

105
Q

clo3-

A

chlorate

106
Q

clo4-

A

perchlorate

107
Q

hpo4 2-

A

hydrogen phosphate

108
Q

c2o4 2-

A

oxalate

109
Q

so3 2-

A

sulfite

110
Q

so4 2-

A

sulfate

111
Q

co3 2-

A

carbonate

112
Q

cro4 2-

A

chromate

113
Q

cr2o7 2-

A

dichromate

114
Q

si03 2-

A

silciate

115
Q

po3 3-

A

phosphite

116
Q

po4 3-

A

phosphate

117
Q

nh4+

A

ammonium

118
Q

2 kinds of binary compounds

A

compounds that contain a metal and nonmetal

119
Q

naming type 1 binary compounds

A

1) the cation is always named first and the anion is second
2) the cation has the same name as the element
3) the anion is named by taking the first part of the element name (root) and adding -ide

120
Q

type 1 binary compound examples

A

nacl: na+, cl-, sodium chloride

ki: k+, i-, potassium iodide

cas: ca2+, s2-, calcium sulfide

121
Q

naming type 2 binary compounds

A

1) type 2 compounds are those in which the metal cation has more than one valence number
2) the basic rule is the same as type 1
3) group 1 and 2 metals are always type 1
4) transition metals are almost always type 2
5) a roman numeral is placed in the name of the compound to indicate the valence of the cation.

122
Q

type 2 binary compound examples

A

fecl: iron (iii) chloride

sn3n3: tin (ii) nitride

(hg2)3p2: mercury (i) phosphide

123
Q

mercury exceptions (again)

A

roman numeral doesn’t indicate the valcene for mercury but rather refers to the subscript. both have a valence of 2+.
- mercury (I) has a subscript of 2
- mercury has a subscript of 1 (invisible)

124
Q

type 3 binary compounds

A

compounds that contain only non metals

125
Q

naming type 3 binary compounds

A

1) the first element in the formula is named first, full element name is used
2) the second element is named as if it’s an anion
3) prefixes are used to denote the number of atoms present
4) the prefix mono- is never used for naming the first element. for example, co is called carbon monoxide and not mono carbon monoxide

126
Q

prefixes

A

mono -> 1
di -> 2
tri -> 3
tetra -> 4
penta -> 5
hexa -> 6
heat -> 7
octa -> 8

127
Q

type 3 binary compounds examples

A

no: nitrogen monoxide

n2o: dinitrogen monoxide

if5: iodine pentaflouride

p4o6: tetraphosphorus hexoxide

128
Q

naming compounds containing polyatomic ions

A

1) put the name of the cation first and the name of the anion second
2) use roman numerals after the name for the cation for type 2 compounds

129
Q

compounds containing polyatomic ions examples

A

ca(oh)2: calcium hydroxide

na3po4: sodium phosphide

(nh4)2 cr2o7: ammonium dichromate

(hg2)3(po4): mercurey (i) phosphate

130
Q

naming acids

A

1) h & something else: use prefix HYDRO + root of anion + ic acid

2) polyatomic ions
- h+ate becomes root + ic
- h+ite becomes root + ous

131
Q

1 mole contains

A

6.022 * 10^23

132
Q

indications of a chemical reaction

A

1) evolution of energy as heat and light
2) production of gas
3) formation of precipitate
4) color change

133
Q

characteristics of chemical reactions

A

1) equation must represent known facts
2) the equation must contain the correct formulas for the reactants and products
3) law of conversation of mass

134
Q

aluminum + iron (III) oxide -> aluminum oxide + iron

A

2al + fe2o3 -> al2o3 + 2fe

135
Q

sodium hydroxide + iron (ii) chloride -> sodium chloride + iron (ii) hydroxide

A

12naoh + fecl2 -> 2nacl fe(oh)2

136
Q

synthesis

A

A+B

137
Q

double replacement

A

AB + CD -> AD + BC

138
Q

single replacement

A

A + BC -> AC + B

139
Q

decomposition

A

AB -> A + B

140
Q

combustion

A

whatever -> CO2 + H2O

141
Q

single replacement rules

A

if the outside atom/molecule is less reactive, then single replacement won’t happen

142
Q

transition metals

A

strong, dense, less reactive, used in jewlery

143
Q

metalloid

A

groups 3-12, properties of metals & non metals, semiconductors

144
Q

inner transition elements

A

bottom 2 rows, many radioactive, man-made

Chemistry H (11 decks)

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