FINALS! Flashcards
(119 cards)
1
Q
homogeneous
A
same properties throughout
2
Q
heterogeneous
A
different properties in different parts of mixture
3
Q
elements
A
cannot be broken down chemically into simpler substance
4
Q
compounds
A
can be broken down chemically into elements
5
Q
precision
A
degree of agreement among several measurements of the same quality
6
Q
accuracy
A
the agreement of a particular value with the true value
7
Q
sigfig zero rules
A
- non-zero always count as sigfigs
a) leading 0s never count
2) captive 0s always count
3) trailing 0s only included if theres decimal points
8
Q
density formula
A
mass/volume
9
Q
atomic number
A
protons/electrons
10
Q
atomic mass
A
protons + neutrons
11
Q
ground state orbitals
A
lowest energy level/orbital
can be moved up by heat, electricity, light (excited)
12
Q
quantum mechanics
A
how small particles behave
13
Q
energy levels
A
measures fixed energy e-
since e- cannot exist between rungs, a quantum is the exact energy needed to move an e- up a rung
14
Q
principal quantum number
A
denotes the energy level e- is located in
max in an energy level: 2n^2
15
Q
aufbau
A
electrons enter the lowest energy level first
16
Q
pauli exclusion
A
2 electrons max per orbital
17
Q
hund’s rule
A
electrons don’t pair up unless they have to
18
Q
alkali metals
A
group 1
most reactive
not found in nature
reacts with air and water
19
Q
alkaline earth metals
A
group 2
reactive, but not as much as alkali
20
Q
transition metals
A
group 3-12
all metals
least reactive on periodic table
found in nature
21
Q
rare earth metals
A
bottom 2 rows
22
Q
lanthanides
A
1st bottom row on periodic table
soft metals, not that rare
23
Q
actinides
A
2nd bottom ro won periodic table
radioactive, synthetic
24
Q
halogens
A
group 17
most reactive nonmetals
25
noble gases
group 18
rarely combine - low reactivity
26
group on periodic table
up & down, column
27
period on periodic table
left & right, row
28
atomic radius trend
1) increases down a group
2) decreases across a period
29
electronegativity trend
1) decreases down a group
2) increases across a period
30
ionization energy trend
1) decreases across a group
2) increases across a period
31
ion size trend
1) larger when anion (gain electron)
2) smaller when cation (lose electron)
32
metallic character trend
1) increases down a group
2) decreases down a period
33
ionic bond
attractions between oppositely charged ions
34
covalent bond
2 nonmetals bonding by sharing electrons
35
non-polar covalent
equal sharing of electrons
diatomic molecules
36
polar covalent
unequal sharing of electrons
electrons spend more time around the nonmetallic tom
charge seperation - dipole movement
37
metallic bonds
electrostatic attraction between cations (2 metals)
38
covalent network solids
combinations of nonmetals
hard and brittle
extreme melting and boiling points
interconnected, insoluble
39
1 central, 2 atoms
linear
40
1 central, 3 atoms
trigonal planar
41
1 central, 4 atoms
tetrahedral
42
1 central + 1 lone pair, 2 atoms
bent
43
1 central + 1 lone pair, 3 atoms
trigonal pyramidal
44
1 central + 2 lone pairs, 2 atoms
bent
45
polar bond vs polar molecule
polar bond: unequal sharing of e-
polar molecule: non symmetrical shape, lone pair on central atom
46
non polar bond vs non polar molecule
nonpolar bond: equal sharing of e-
nonpolar molecule: symmetrical molecular shape
47
hydrogen bonds
dipole-dipole
strong intermolecular froce
occurring between hydrogen atoms with fluorine, oxygen, nitrogen
48
type 1 binary compound
metal present forms 1 type of cation
1) cation first, anion 2nd
2) cation has same name elemtn
3) anion + root ide
49
type 2 binary compound
metal present forms 2+ cations with diff charges
1) metal cation has more than 1 valence number
2) group 1&2 always type 2
3) transition metals almost always type 2
4) roman numeral placed to indicate valence
50
mercury exception (type 2 binary compound)
roman numeral refers to subscript b/c mercury I and mercury II both have valence of 2+
51
type 3 binary compound
2 nonmetals
1) named with full element
2) prefixes denote # of atoms present
3) mono never used for first element
mono:1
di:2
tri:3
tetra:4
penta:5
hexa:6
hepta:7
octa:8
52
empirical formula
lowest whole number ratio of atoms in a compound
53
molecular formula
the true number of atoms of each elements in the formula
54
empirical & formula proccess
1) get the % of eachh part
2) divide by molar mass
3) divide by smallest
4) round
55
A + X -> AX
synthesis/combination
56
AX -> A + X
decomposition
57
A + BX -> AX + B
single replacement reaction
58
AX + BY -> AY + BX
double replacement
59
CxHy + O2 -> CO2 + H2O
combustion
substance + oxygen --> energy (light, heat)
60
evaporation
when liquid converts to gas when the liquid is NOT boiling (on the surface)
61
boiling
a conversion of a liquid to a gas or vapor through the whole substance
62
vaporization
conversion of ANY liquid molecule into a gas molecule
63
q = m * cp * ∆T
q = thermal energy
m = mass (g)
cp = specific heat
∆T = change in temp
64
cp of h2o
4.184 J/gºK
65
heating curves, increasing sections
q = m * cp * ∆T
endothermic
66
heating curves, plateaus
1) melting: q = mol * ∆Hfus
2) boiling q = mol * ∆Hvap
67
kinetic theory of gases
1) gases are mostly empty space - no forces of attraction or repulsion
2) gases are in constant motion
3) collisions between gas particles are perfectly elastic
68
boyles law
inverse
p1v1=p2v2
69
charles law
proportional
v1/t1 = v2/t2
70
gay-lussac law
proportional
p1/t1=p2/t2
71
ideal gas law
PV=nRT
V = L
n = moles
R = gas constant
T = K = C+273
72
stp
1 atm = 101.3 kPA = 760 mmHg
73
dalton's law of partial particles
p total = p1 + p2 + p3
74
diffusion
effusion: passage of gas particles through a small opening
diffusion: the moment of particles for regions of high concentration to low concentration
75
grahams law
rate1/rate2 = sqrt (m2/m1)
76
ideal gas
- no volume, no attraction/repulsion, all collisions are elastic
don't exist
closest at low pressure and high temperature.
77
properties of water
- universal solvent
- bent, v-shaped, 105
- polar bonds (covalent between o-h)
- oxygen slightly negative, hydrogen slightly positive
- water molecules cant form h-bonds with air molecules, only attracted to h-bonds in the body of the liquid
78
colloids
milky and cloudy
79
surface tension
molecules @ top are only pulled to inside
molecules in middle are attracted in all directions
causes droplets to minimize surface area
80
degree of solubility
1) nature of solute & solvent "like dissolves like"
2) temperature: increase temp increase solubility
3) pressure (FOR GASES) increase pressure, increase solubility
81
rate of solution
1) increase TEMP, dissolve faster b/c kinetic energy
2) smaller PARTICLES dissolve faster
3) STIRRING dissolves faster bc concentration gradient
4) already dissolved solute dissolves slower bc less concentration gradient
82
gas in liquid solubility
1) temperature: increase temp, less gas dissolved
2) pressure: increase pressure, more gas dissolved
(solids more soluble as temp increases / gas less soluble as temp increases)
83
solubility curves
unsatured: more solute dissolves
saturated: no more solute dissolves
supersaturated: unstable, crystals form
84
molarity
moles of solute/liters of solution
85
molality
moles of solute/kg of solvent
86
dilutions
made by adding more solvent to a solution
moles o/solute before dilution = moles o/solute after dilution
only concentration changes
m1v1=m2v2
87
colligative properties
boiling point elevation
freezing point depression
88
boiling point elevation
∆tb = i * kb * m
∆Tb = change in boiling point
I = # of subatomic particles
kb = molal boiling point constant (0.512 for water)
m = molality
89
freezing point depression
∆Tf = i * kf * m
∆Tf = change in freezing pt
i = # of subatomic particles
Kf = molal freezing point constant (1.86 for water)
m - molality
90
potential energy
stored energy
91
kinetic energy
energy of motion
92
temperature vs heat
temp: measure of the average kinetic energy of random motions of particles in substances
heat: measure of the total amount of energy
93
specific heat capacity
amount of heat needed to increase temp of 1g of a substance by 1c
q = m * cp * ∆T
q = joules
m = mass
cp = specific heat
∆t = change in temp
94
enthalpy
heat of fusion: energy needs to melt one mole
heat of vaporization: energy needed to boil one mole
95
collision theory
1) frequency of collisions: high # of collisions needed for reactions to occur
2) effectiveness of collisions: particles collide at proper angles & enough energy
96
rates of reaction
1) inc temp, inc rate
2) inc concentration, inc rates
3) inc pressure, inc rate
4) dec particle sizxe
5) cataylsts
97
la chatelier's pirnciple
if stress is applied to a system at equilibrium, the system changes to relieve the stress to establish a new equilibrium
98
stress
1) change in concentration
2) change in temp
3) change in pressure
99
stress
A + B <-> C + D
increase A or B shifts right (makes more products)
decrease A or B shifts left (makes more reactants)
100
endo/exo stress
A+B <-> C+D
A+B <-> C+D + heat
exothermic, shifts right
A+B+heat <-> C+D
endothermic, shifts left
101
Keq
equilibrium constant
[products]/[reactants]
([C]^c * [D]^d)/([A]^a * [B]^b)
exclude solids and pure liquids
102
keq><=1
Keq>1, products favored
Keq=1, neither favored
Keq<1, reactants favored
103
acid basic properties
changes litmus red
produces H+ when dissolved in water
104
naming acids
1) -ide → starts with hydro, suffix ic, end acid
2) -ite → suffix ous, end acid
3) -ate → suffix ic, end acid
105
base basic properties
changes litmus blue
produces OH- when dissolved in water
106
bronsted-lowry acids
acid: H+ donor (proton donor)
base: H+ acceptor (proton acceptor)
HBr+H2O <-> H3O+ + Br-
acid. base.
107
conjugate acid-base paris
2 substances that differ by 1 H+
acid -> conjugate base
H2O -> OH-
base -> conjugate acid
NH3 -> NH4+
108
acid strenght
strong acids completely dissociate in water
HCl -> H+ + Cl-
weak acids only partially ionize in ater
H3COOH (aq) <-> CH3COO-(aq) + H+
109
neutralizing reactions
all neutralization runs are double replacement
salt is an ionic compound formed form an acid (anion) and base (cation)
110
normality
normality (N) is a unit of concentration in acid-base titrations
N1V1=N2V2
acids
monoprotic (N) = (M)
diprotic N=2M
triprotic N=3M
bases
1 OH- ion N=M
2 OH- ion N=2M
3 OH- ion N=3M
111
buffers
solutions that resist changes in pH
the buffer cannot control the pH when too much acid/base is added
112
titrations
a method to determined the concentration of a solution using neutralizing reactions
113
titration indicators
indicators: added to signal when neutralization has occured
changes color @ end point
neutralization @ equivalence point
114
titration curves
1) weak acids neutralized by strong bases produce basic salt solutions
2) strong acids neutralized by weak bases produce acidic salt solutions
3) strong acids neutralized by strong bases produce neutral salt solutions
115
oxidation
loss of electrons (atoms becomes more positive)
116
reduction
gain of electrons (atoms became more negative)
117
oxidizing agent
causes oxidation of another element, gets reduced
118
reducing agent
causes reduction of another element, gets oxidized
119
redox
reactions with reduction and oxidation
synthesis, decomposition, single reaplcement
NOT double replacement, neutralization
