MIDTERM! Flashcards
(144 cards)
1
Q
mass vs weight
A
- mass is a measure of object resistance
- weight is measured through earth’s gravity
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
3
Q
random error
A
the measured value has an equal chance of being too high or too low (not precise, not accurate)
4
Q
systematic error
A
the measured value will always high (or low)/ dates off by the same amount each time (precise, not accurate)
5
Q
sig figs non zero integers
A
always count as significant figures
6
Q
sig fig zero rules
A
- leading zeros - never count (0.0025 is only 2 sig figs)
- captive zeros - always count (1.008 is 4 sig figs)
- trailing zeros - only count if # has a decimal point
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
8
Q
regular notation
A
standard way to weight numbers
9
Q
scientific notation
A
short handed way of writing (2.8 *10^24315)
10
Q
giga (G)
A
10^9
11
Q
mega (M)
A
10^6
12
Q
kilo (k)
A
10^3
13
Q
deci (d)
A
10^-1
14
Q
centi (c)
A
10^-2
15
Q
milli (m)
A
10&-3
16
Q
micro (u)
A
10^-6
17
Q
nano (n)
A
10^-9
18
Q
pico (p)
A
10^-12
19
Q
femto (f)
A
10^-15
20
Q
C = (F-32) * 5/9
A
fahrenheit to celsius
21
Q
celsius to fahrenheit
A
f = 9/5c + 32
22
Q
celsius to kelvin
A
k = c + 273
23
Q
matter classifies into
A
- pure substances
- mixtures
24
Q
pure substances classifies into
A
- elements (one type of atom)
- compound (more than one type of atoms)
25
mixtures classifies into
1. homogeneous (uniform)
2. heterogenous (non uniform)
26
proton basic info
+ charge, 1 amu
27
neutron charge
0 charge, 1 amu
28
electron basic info
- charge, 1/1836 amu
29
isotopes
atoms of the same elements with different #s of neutrons (mass changes)
30
ions
charged atom (gain or lose e-) (mass doesn't change)
31
atom box LABELS
1. top right, mass number
2. top left, charge
3. bottom left, atomic #
32
weighted average
((relative abundance * mass of isotope) + (relative abundance * mass of isotope 2))/100
33
proton and electron relation
equal, unless there's a charge
34
proton and neutron relation
mass # - neutrons = protons
nass # - protons = neutrons
35
aristotle
340 BCE
- greek philosopher aristotle thought fire, water, air, and earth were the building blocks of everything
36
dalton
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
37
rutherford
- scattering experiment, rutherford sent alpha particles through a thin sheet of gold
- allowed him to discover nucleus
38
bohr
- 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
39
jj thompson
- discovered the electron
40
chadwick
- discovered atoms not only consist of protons and electrons but also neutrons!
- neutral subatomic particle has around the same mass as a proton
41
democritus
- 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"
42
schrodinger & heisenberg
- 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)
43
chemical properties
describe a substances ability to change to a difference substance
44
element
cannot be broken down chemically into simpler substance
45
compounds
can be broken down chemically into elements (are 2 or more atoms bounded together)
46
solvent
does the dissolving
47
solute
gets dissolved
48
alloy
one or more slides dissolved in another solid
49
lowest energy level
where electron starts from is ground state
- electron configuration written in lowest energy
- atomic spectra
50
excited state
- heat, electricity, or light can move up to different energy levels
- when it falls back to ground state, it gives back energy as light
51
returning to ground state
- 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
52
quantum mechanics
an explanation of how small particles behave - an explanation for subatomic particles and atoms as waves
53
classical mechanics
describes the motions of bodies much larger than atoms
54
heisenberg uncertainty principles
you cannot know both the position and momentum of an electron (where its going vs where it is)
- warner heisenberg!
55
ernest rutherford model
- 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
56
niel bohr's model
- 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)
57
ladder rungs
- 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
58
quantum mechanical medol
- 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
principle quantum number (n)
- 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
atomic orbitals
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
sublevels
s, p, d, f
62
s (spherical)
1 orbital, 2 maximum electrons
63
p (dumbell)
3 orbitals, 6 maximum electrons
64
d (clover leaf)
5 orbitals, 10 maximum electrons
65
f (complicated)
7 orbitals, 14 maximum electrons
66
energy level 1!!
- only s sublevel
- 1s2 (1 orbital) w/ only 2e-
2 total e-
67
energy level 2!!
- has s and p sublevels
- 2s2 (2 orbitals) w 2e-
- 2p6 (3 orbitals) w 6e-
2s2 2p6 ~ 8 total e-
68
energy level 3!!
- 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
energy level 4!!
- 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
aufbau principle
electrons enter the lowest energy first (causes difficulties because overlap of orbitals of different energies)
71
pauli exclusion principle
2 electrons max per orbital ~ different spins. no two electrons in an atom have the same 4 quantum numbers
72
hund's rule
when electrons occupy orbitals of equal energy, they don't air up until they have to
73
chromium exception
4s2 3d4 --> 4s1 3d5
74
copper exception
4s2 3d9 --> 4s1 3d10
75
group number
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
atomic size
- increases going down a group
- decreases going across a period
77
ion size
- cations smaller than the atoms they come from
- anions larger than the atoms they come from
78
ionization energy
- 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
electronegativity
- 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
metallic character
a measure of how easily an atom uses an e-
- non metallic at F, CL
- most metallic at Cs, Fr
81
alkali metals
tend to form +1 ions
82
alkali earth metals
tend to form +2 ions
83
halogens
tend to form -1 ions
84
group 1 (+1)
li
na
k
pb
cs
85
group 2 (+2)
be
mg
ca
sr
ba
86
group 3-12
nothing lol
87
group 13 (+3)
al
88
group 14
nothing lol
89
group 15 (-3)
n
p
as
90
group 16 (-2)
o
s
se
91
group 17 (-1)
f
cl
br
92
group 18 (0)
nothing lol
93
h2po4-
dihydrogen phosphate
94
c2h3o2
acetate
95
hso3-
hydrogen sulfite
96
hso4-
hydrogen sulfate (bisulfate)
97
hco3-
hydrogen carbonate (bicarbonate)
98
no2-
nitrite
99
no3-
nitrate
100
cn-
cyanide
101
oh-
hydroxide
102
mno4-
permanganate
103
clo-
hypochlorite
104
clo2-
chlorite
105
clo3-
chlorate
106
clo4-
perchlorate
107
hpo4 2-
hydrogen phosphate
108
c2o4 2-
oxalate
109
so3 2-
sulfite
110
so4 2-
sulfate
111
co3 2-
carbonate
112
cro4 2-
chromate
113
cr2o7 2-
dichromate
114
si03 2-
silciate
115
po3 3-
phosphite
116
po4 3-
phosphate
117
nh4+
ammonium
118
2 kinds of binary compounds
compounds that contain a metal and nonmetal
119
naming type 1 binary compounds
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
type 1 binary compound examples
nacl: na+, cl-, sodium chloride
ki: k+, i-, potassium iodide
cas: ca2+, s2-, calcium sulfide
121
naming type 2 binary compounds
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
type 2 binary compound examples
fecl: iron (iii) chloride
sn3n3: tin (ii) nitride
(hg2)3p2: mercury (i) phosphide
123
mercury exceptions (again)
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
type 3 binary compounds
compounds that contain only non metals
125
naming type 3 binary compounds
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
prefixes
mono -> 1
di -> 2
tri -> 3
tetra -> 4
penta -> 5
hexa -> 6
heat -> 7
octa -> 8
127
type 3 binary compounds examples
no: nitrogen monoxide
n2o: dinitrogen monoxide
if5: iodine pentaflouride
p4o6: tetraphosphorus hexoxide
128
naming compounds containing polyatomic ions
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
compounds containing polyatomic ions examples
ca(oh)2: calcium hydroxide
na3po4: sodium phosphide
(nh4)2 cr2o7: ammonium dichromate
(hg2)3(po4): mercurey (i) phosphate
130
naming acids
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
1 mole contains
6.022 * 10^23
132
indications of a chemical reaction
1) evolution of energy as heat and light
2) production of gas
3) formation of precipitate
4) color change
133
characteristics of chemical reactions
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
aluminum + iron (III) oxide -> aluminum oxide + iron
2al + fe2o3 -> al2o3 + 2fe
135
sodium hydroxide + iron (ii) chloride -> sodium chloride + iron (ii) hydroxide
12naoh + fecl2 -> 2nacl fe(oh)2
136
synthesis
A+B
137
double replacement
AB + CD -> AD + BC
138
single replacement
A + BC -> AC + B
139
decomposition
AB -> A + B
140
combustion
whatever -> CO2 + H2O
141
single replacement rules
if the outside atom/molecule is less reactive, then single replacement won't happen
142
transition metals
strong, dense, less reactive, used in jewlery
143
metalloid
groups 3-12, properties of metals & non metals, semiconductors
144
inner transition elements
bottom 2 rows, many radioactive, man-made
