FINALS! Flashcards
homogeneous
same properties throughout
heterogeneous
different properties in different parts of mixture
elements
cannot be broken down chemically into simpler substance
compounds
can be broken down chemically into elements
precision
degree of agreement among several measurements of the same quality
accuracy
the agreement of a particular value with the true value
sigfig zero rules
- 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
density formula
mass/volume
atomic number
protons/electrons
atomic mass
protons + neutrons
ground state orbitals
lowest energy level/orbital
can be moved up by heat, electricity, light (excited)
quantum mechanics
how small particles behave
energy levels
measures fixed energy e-
since e- cannot exist between rungs, a quantum is the exact energy needed to move an e- up a rung
principal quantum number
denotes the energy level e- is located in
max in an energy level: 2n^2
aufbau
electrons enter the lowest energy level first
pauli exclusion
2 electrons max per orbital
hund’s rule
electrons don’t pair up unless they have to
alkali metals
group 1
most reactive
not found in nature
reacts with air and water
alkaline earth metals
group 2
reactive, but not as much as alkali
transition metals
group 3-12
all metals
least reactive on periodic table
found in nature
rare earth metals
bottom 2 rows
lanthanides
1st bottom row on periodic table
soft metals, not that rare
actinides
2nd bottom ro won periodic table
radioactive, synthetic
halogens
group 17
most reactive nonmetals
noble gases
group 18
rarely combine - low reactivity
group on periodic table
up & down, column
period on periodic table
left & right, row
atomic radius trend
1) increases down a group
2) decreases across a period
electronegativity trend
1) decreases down a group
2) increases across a period
ionization energy trend
1) decreases across a group
2) increases across a period
ion size trend
1) larger when anion (gain electron)
2) smaller when cation (lose electron)
metallic character trend
1) increases down a group
2) decreases down a period
ionic bond
attractions between oppositely charged ions
covalent bond
2 nonmetals bonding by sharing electrons
non-polar covalent
equal sharing of electrons
diatomic molecules
polar covalent
unequal sharing of electrons
electrons spend more time around the nonmetallic tom
charge seperation - dipole movement
metallic bonds
electrostatic attraction between cations (2 metals)
covalent network solids
combinations of nonmetals
hard and brittle
extreme melting and boiling points
interconnected, insoluble
1 central, 2 atoms
linear
1 central, 3 atoms
trigonal planar
1 central, 4 atoms
tetrahedral
1 central + 1 lone pair, 2 atoms
bent
1 central + 1 lone pair, 3 atoms
trigonal pyramidal
1 central + 2 lone pairs, 2 atoms
bent
polar bond vs polar molecule
polar bond: unequal sharing of e-
polar molecule: non symmetrical shape, lone pair on central atom
non polar bond vs non polar molecule
nonpolar bond: equal sharing of e-
nonpolar molecule: symmetrical molecular shape
hydrogen bonds
dipole-dipole
strong intermolecular froce
occurring between hydrogen atoms with fluorine, oxygen, nitrogen
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
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
mercury exception (type 2 binary compound)
roman numeral refers to subscript b/c mercury I and mercury II both have valence of 2+
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
empirical formula
lowest whole number ratio of atoms in a compound
molecular formula
the true number of atoms of each elements in the formula
empirical & formula proccess
1) get the % of eachh part
2) divide by molar mass
3) divide by smallest
4) round
A + X -> AX
synthesis/combination
AX -> A + X
decomposition
A + BX -> AX + B
single replacement reaction
AX + BY -> AY + BX
double replacement
CxHy + O2 -> CO2 + H2O
combustion
substance + oxygen –> energy (light, heat)
evaporation
when liquid converts to gas when the liquid is NOT boiling (on the surface)
boiling
a conversion of a liquid to a gas or vapor through the whole substance
vaporization
conversion of ANY liquid molecule into a gas molecule
q = m * cp * ∆T
q = thermal energy
m = mass (g)
cp = specific heat
∆T = change in temp
cp of h2o
4.184 J/gºK
heating curves, increasing sections
q = m * cp * ∆T
endothermic
heating curves, plateaus
1) melting: q = mol * ∆Hfus
2) boiling q = mol * ∆Hvap
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
boyles law
inverse
p1v1=p2v2
charles law
proportional
v1/t1 = v2/t2
gay-lussac law
proportional
p1/t1=p2/t2
ideal gas law
PV=nRT
V = L
n = moles
R = gas constant
T = K = C+273
stp
1 atm = 101.3 kPA = 760 mmHg
dalton’s law of partial particles
p total = p1 + p2 + p3
diffusion
effusion: passage of gas particles through a small opening
diffusion: the moment of particles for regions of high concentration to low concentration
grahams law
rate1/rate2 = sqrt (m2/m1)
ideal gas
- no volume, no attraction/repulsion, all collisions are elastic
don’t exist
closest at low pressure and high temperature.
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
colloids
milky and cloudy
surface tension
molecules @ top are only pulled to inside
molecules in middle are attracted in all directions
causes droplets to minimize surface area
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
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
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)
solubility curves
unsatured: more solute dissolves
saturated: no more solute dissolves
supersaturated: unstable, crystals form
molarity
moles of solute/liters of solution
molality
moles of solute/kg of solvent
dilutions
made by adding more solvent to a solution
moles o/solute before dilution = moles o/solute after dilution
only concentration changes
m1v1=m2v2
colligative properties
boiling point elevation
freezing point depression
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
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
potential energy
stored energy
kinetic energy
energy of motion
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
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
enthalpy
heat of fusion: energy needs to melt one mole
heat of vaporization: energy needed to boil one mole
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
rates of reaction
1) inc temp, inc rate
2) inc concentration, inc rates
3) inc pressure, inc rate
4) dec particle sizxe
5) cataylsts
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
stress
1) change in concentration
2) change in temp
3) change in pressure
stress
A + B <-> C + D
increase A or B shifts right (makes more products)
decrease A or B shifts left (makes more reactants)
endo/exo stress
A+B <-> C+D
A+B <-> C+D + heat
exothermic, shifts right
A+B+heat <-> C+D
endothermic, shifts left
Keq
equilibrium constant
[products]/[reactants]
([C]^c * [D]^d)/([A]^a * [B]^b)
exclude solids and pure liquids
keq><=1
Keq>1, products favored
Keq=1, neither favored
Keq<1, reactants favored
acid basic properties
changes litmus red
produces H+ when dissolved in water
naming acids
1) -ide → starts with hydro, suffix ic, end acid
2) -ite → suffix ous, end acid
3) -ate → suffix ic, end acid
base basic properties
changes litmus blue
produces OH- when dissolved in water
bronsted-lowry acids
acid: H+ donor (proton donor)
base: H+ acceptor (proton acceptor)
HBr+H2O <-> H3O+ + Br-
acid. base.
conjugate acid-base paris
2 substances that differ by 1 H+
acid -> conjugate base
H2O -> OH-
base -> conjugate acid
NH3 -> NH4+
acid strenght
strong acids completely dissociate in water
HCl -> H+ + Cl-
weak acids only partially ionize in ater
H3COOH (aq) <-> CH3COO-(aq) + H+
neutralizing reactions
all neutralization runs are double replacement
salt is an ionic compound formed form an acid (anion) and base (cation)
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
buffers
solutions that resist changes in pH
the buffer cannot control the pH when too much acid/base is added
titrations
a method to determined the concentration of a solution using neutralizing reactions
titration indicators
indicators: added to signal when neutralization has occured
changes color @ end point
neutralization @ equivalence point
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
oxidation
loss of electrons (atoms becomes more positive)
reduction
gain of electrons (atoms became more negative)
oxidizing agent
causes oxidation of another element, gets reduced
reducing agent
causes reduction of another element, gets oxidized
redox
reactions with reduction and oxidation
synthesis, decomposition, single reaplcement
NOT double replacement, neutralization
