unit 2 Flashcards
includes all unit 2 power point information and practice problems (not including math)
periodic law original
when the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically
quantum mechanics explains
why periodic trends exist
metals preperties
solids at room temp
lustrous
good conductors
malleable
ductile
loses electrons to form cations
nonmetals properties
gases/brittle solids at room temp
poor conductors
gains electrons to form anions
metalloids
poor conductors at room temp
better conductors at high temps
alkali metals properties
oxides dissolve in water to form strong basic solutions
corrode in air to dull gray appearance
react vigorously to produce hydrogen
readily form compounds with nonmetals
s1 electron configuration
cations with 1+charge
gets larger as move down group
alkaline earth metals
not as reactive as alkali metals
readily form compounds with nonmetals
forms alkaline oxides, some have low solubility in water
s2 electron configuration
cations with +2 charge
halogens properties
most reactive nonmetals
F and Cl are gases, Br is liquid, I is solid
exist as diatomic molecules
readily forms compounds with metals and nonmetals
s2p5 electron configuration
gains electrons to have a 1- charge
shares electrons with nonmetals
noble gases
unreactive
s2p6 outer electron configuration
stable octet
low reactivity
isoelectronic
has same electron configuation
transition metals properties
closer to nucleus than representative metals
solid
high melting and boiling points
similarities in properties in a group and across a period
forms cations with multiple charges, forms complex ions
some have characteristic colors
3 factors of periodic trends
nuclear charge
number of core electrons
distance of outer electrons from the nucleus
coulomb’s law
e = 1/4piCo q1q2/r
e: potential energy
q1q2: charges
r: separation
coulomb’s law
strength of interaction increases as charge increases
attractive force between electron and nucleus depends on magnitude of nuclear charge (atomic number) and average distance between nucleus and electron (energy levels)
shielding
repulsions cause electron to have net reduced attraction to nucleus
effective nuclear charge
amount of attraction electron feels for nucleus
2 types of shielding
shielding by core electrons (efficient)
shielding from outer electrons (not efficient)
Zeff =
Z (actual charge of nucleus) - S (charge screened by other electrons)
atomic radius
1/2 distance between 2 atoms
bond lengths calculation
add atomic radii or 2 bonding atoms
cations are _ than their neutral atoms
smaller
anions are _ than their original atom
larger
ionization energy
minimum energy needed to remove an electron
greater IE = harder to remove an electron
increases are electrons are removed, significant jump when an entire energy level is removed
first ionixation energy
energy needed to remove the first electron
lowest
electron affinity
energy change when atom gains electron
greater attraction between atom + added electron = more negative EA
positive EA = electron won’t attach itself
halogens have most negative
electronegativity
ability of an atom in a molecule to attract shared electrons to itself
higher EN = pulls bonded electron closer
smaller atoms = higher EN
spectroscopy
studies transmission/reflection of different frequencies of electromagnetic spectrum by a sample of matter
photoelectron spectroscopy
core more difficult to remove than valence
uses higher energy light: UV to X-rays
UV studies valence
X-rays studies core
how to do photoelectron spectroscopy
- beam of x-rays shone on sample (e=hv)
- atoms of sample release electrons from any energy level (photoelectric effect)
- kinetic energy measured (binding energy = hv - KE)
PES spectrum
graph of binding energy (x-axis) and relative number of electrons (y-axis)
height of each peak is proportional to # of electrons of equal energy
UV and visible light spectroscopy
white visible light shone on colorful molecules + transition metal ions
doesn’t remove electrons from atom
excites electrons to higher energy state
electrons absorb certain wavelengths and transmit other wavelengths, producing color
reference blank
when other compounds in a solution absorbs the same wavelengths as compound being analyzed
beer’s law
A = Ebc
A=absorbance
E=molar absorbance
b=path length
c=concentration
calibration curves
direct relationship between absorbance and concentration
prepare solutions of known concentrations and analyze at ymax
plot absorbance as function of concentration
used to find the concentration of an unknown solution
meter
m
length
kilogram
kg
mass
second
s
time
Pascal
Pa
pressure
kelvin
K
temp
liter
L
volume
mole
mol
amount of substance
SI system
dominant system of measurement
giga
1,000,000,000
mega
1,000,000
kilo
1,000
hecto
100
deka
10
deci
0.1
centi
0.01
milli
0.001
micro
0.000 001
nano
0.000 000 001
pico
0.000 000 000 001
precision
exactness of measurement
error =
|measured - accepted|/accepted
how did Mendeleev organize the PT?
ordered by increasing atomic mass
elements with similar properties were in the same column
issues with periodic law
doesn’t explain why the column pattern exists
How did Mendeleev’s periodic table differ from the modern periodic table and how did this affect the wording of the periodic law?
Mendeleev: when elements are arranged in order of atomic mass, periodic properties recur
Modern PT is arranged by atomic number (number of protons)
Why do elements in the same group have similar chemical properties?
They have the same number of valence electrons
Which two families contain the most reactive elements? Can you suggest a possible reason for this given their locations in the periodic table?
Alkali metals and halogens
Alkali metals lose their outer s1 electron to gain a noble gas configuration
Halogens add an electron to their outer s2 p5 configuration to gain a noble gas configuration
periodic trends are due to:
strength of attraction between the nucleus and outer electrons
3 factors strength of attraction depends on
nuclear charge
number of core electrons
distance of outer electrons from nucleus
covalent radius
1/2 distance between 2 covalently bonded atoms
metallic radius
1/2 distance between 2 adjacent metal atoms in a metallic crystal
nonbonding radius (van der Waal’s radius)
1/2 distance between 2 non-bonded atoms that are touching
metallic periodic trend
increases across a period and down a group
atomic radius trend
increases down a group
decreases across a period
to calculate bond lengths:
add atomic radii of two bonding atoms
What attractive force is responsible for holding the cloud of electrons in place in atoms?
Nuclear attraction
What effect would a strengthening of nuclear attraaction have on the size of atoms?
Increasing nuclear attraction makes the atom smaller
What might cause a strengthening of that force?
Increasing the nuclear charge (adding protons)
What might contribute to a weakening of that force?
Increasing the shielding effect by adding core electrons
Why does the quantum mechanical model description of multi-electron atoms make it difficult to define a precise atomic radius?
Atoms do not have precise boundaries. Orbitals represent the 90 % probability of finding an
electron in a volume of space
When nonbonded atoms come up against one another, what determines how closely the nuclear centers can approach?
Repulsion between each of the atom’s outermost electrons
Which of the three factors that contribute to atomic size, predominates as we move across a period? What is the result on atomic size?
nuclear attraction predominates (no new energy levels added) making the atom smaller
Which of the three factors contribute to atomic size predominates as we move down a group? What is the result on atomic size?
distance of outer electrons from the nucleus and shielding both increase making the atom larger
what happens after cations form from neutral atoms?
Electrons are removed from energy levels furthest from nucleus
Electron-electron repulsions decrease
what happens when an anion forms from its neutral atom?
Electrons are added to energy levels further from the nucleus
Electron-electron repulsions increase
isoelectronic series
group of ions with the same number of electrons
ionic radius periodic trend
increases down a group
decreases across a period
first ionization energy
energy needed to remove the first electron
ionization energy _ when successive electrons are removed
increases
ionization energy trend
increases across a period
decreases down a group
electron affinity periodic trend
decreases down a group
electronegativity trend
increases across a a period
decreases down a group
effective nuclear charge (Zeff) periodic trend
increases across a a period
constant down a group
distance from nucleus / shielding periodic trend
constant across a period
increases down a group
t/f: Cations are larger than their corresponding neutral atoms
False: cations are smaller than their neutral atoms
t/f: Li1+ is smaller than Li
True: cations are smaller than their neutral atoms
t/f: Cl1- is bigger than I1-
False: ionic radius increases down a group for ions with the same charge
For isoelectronic ions, how are effective nuclear charge and ionic radius related?
They are inversely related-as Zeff increases, ionic radius decreases
why is O2- is larger than O
Because O2- has gained two more electrons, electron-electron repulsion is increased making the ion bigger than the neutral atom
why is S2- is larger than O2
ionic radius increases down a group
why is S2- is larger than K1+
These are an isoelectronic series (3s2 3p6) so increasing atomic number=increasing nuclear charge. S2- has a smaller nuclear charge than K1+ making S2- larger
why is K1+ is larger than Ca2+
These are an isoelectronic series (3s2 3p6) so increasing atomic number=increasing nuclear charge. K1+- has a smaller nuclear charge than Ca2+ making K1+- larger
What is the general relationship between the size of an atom and its first ionization energy?
The larger the atom the lower its first IE
Explain the difference between ionization energy and electron affinity
Ionization energy is the energy change that is required to remove an electron from an atom to form a cation and is usually positive
Electron affinity is the energy change that is required for an atom to gain an electron to form an anion and is usually negative
in PES, the number of peaks determines:
the number of sublevels in an element
smallest peaks = s sublevels
etc.
highest energy peaks appear farther left (orbitals closest to nucleus)
electromagnetic spectrum wavelengths in increasing frequency order
radio
microwave
infrared
visible
ultraviolet
x-ray
gamma ray
reference blank
contains everything that is in the solution that you are trying to measure EXCEPT the thing you are trying to measure
precision and uncertainty relation
lower uncertainty = higher precision
most precise measuring tools
1st: biuret
2nd: measuring cylinder
are non-zero digits significant?
yes
are zero’s between non-zero digits significant?
yes
are leading zero’s significant?
no
are trailing zero’s significant?
only if there is a decimal
multiplication and division sign fig rules
use the same number of sig figs in the answer as the number with the least amount of sig figs
addition and subtraction sig fig rules
the result can only have as many decimal places as the number having the least number of decimal places
One thousand of the base units for pressure
kilopascals (kPa)
One tenth of the base unit for the amount of a substance
decimole (dmol)
One thousandth of the base unit for volume
milliliter (mL)
1 / (1 x 10^6) of the base unit for time
microsecond (us)
1 x 10^-1 of the base unit for length
decimeter (dm)
