Exam 1 Flashcards
Lectures 1-16
1
Q
isotopes
A
atoms of an element with the same number of protons but different number of neutrons
isotopes of same element have similar chemical properities, form same compounds, and display similar reactivities
2
Q
atomic mass
A
- mass of an atom in atomic mass units (amu)
- on periodic table, it is the average mass of the naturally occuring mixture of isotopes
number on periodic table below the element symbol
3
Q
average atomic mass
A
weighted average of all the naturally occuring isotopes of the element
4
Q
atomic number
A
protons
number on periodic table above the element symbol
5
Q
mass number
A
not on periodic table
protons + neutrons
6
Q
quantization of energy
A
stairs analogy
* electron can be on any step but not in between
* higher the energy, higher level = excited state
7
Q
quantum model of atom
A
there is no particular distance at which the electron is located from the nucleus b/c they have wave and particle characteristics
- electrons don’t occupe just a single point in space
- wave behavior of electron is wave function
8
Q
electron density
A
probability of finding an electron in a certain area of space
- results in electron moving rapidly in the orbital b/c faster = more likley to be in certain regions
- proportional to wave function squared
9
Q
orbital
A
locations in space around nucleus at which the probably of finding the electron has higher values
10
Q
Heisenberg Uncertainty Principle
A
it is impossible to know simultaneously the momentum (p) and position of a particle (x)
momentum refers to the particle’s motion/how fast it’s going
p=mv
11
Q
principle quantum number (n)
which shell has the most energy
A
describles electron shell and specifies size of the shell
outer shell = most energy b/c it’s the furthest from the nucleus so more freedom to move
12
Q
angular moment quantum number (l)
A
specifies shape of orbital
l ≤ n-1
- s can hold up to 2 e- and has 1 orbital
- p can hold up to 6 e- and has 3 orbitals
- d can hold up to 10 e- and has 5 orbitals
- f can hold up to 14 e- and has 7 orbitals
- l=0 → s
- l=1 → p
- l=2 → d
- l=3 → f
13
Q
magnetic quantum number (ml)
A
specifies position of orbital
-l ≤ ml ≤ l
ex. for the p-orbital, there is
* ml = -1 for the px orbital
* ml = 0 for the py orbital
* ml = +1 for the pz orbital
14
Q
electron spin (ms)
A
ms = +1/2, -1/2
15
Q
Pauli Exclusion Principle
A
seats in a football stadium analogy
no 2 electrons in an atom can have the same 4 quantum numbers (n, l, ml, ms)
16
Q
number of electrons in a shell formula
A
2(n)^2
17
Q
energy level for orbital diagrams
A
lower = closer to nucleus = lower energy (b/c electrons are closer to protons which means more stable) more stable = less energy
18
Q
shielding
A
inner electrons block outer electrons from protons
- electrons in the same energy level shield each other
19
Q
Aufbau Process
A
- electrons occupy the lowest energy possible
- Hund’s Rule: every orbital is singly occupied w/ one electron before any are doubly occupied
- Pauli exclusion principle: no 2 electrons in an atom have the same 4 quantum numbers
20
Q
Aufbau process exceptions
A
transition metals
lanthanides
actinides
21
Q
how do you know which could represent the electron in the highest energy shell
A
the coefficient
22
Q
inner core electrons
A
those an atom has in common w/ the previous noble gas
23
Q
outer electrons
A
those in the highest energy level (highest n-value)
24
Q
valence electrons
A
involved in bonding
- main group: valence electrons are the outer electrons
- transitional metals: valence electrons are outer electrons and (n-1)d electrons (only if d is not fully filled)
25
# definition and pattern
nuclear charge (Zeff)
net positive charge experienced by an electron (how strongly the nucleus attracts its electrons)
left to right: increases
- more protons
- shielding is relatively constant
bottom to top: increases
- less shells so less shielding
## Footnote
* Nucleus Charge: The nucleus has a positive charge due to protons.
* Electron Shielding: Electrons in the inner shells can "shield" or block some of this positive charge from reaching the outer electrons.
* Net Effect: The effective nuclear charge is the actual positive charge that an outer electron feels, which is less than the total charge of the nucleus due to this shielding effect.
26
# definition and pattern
atomic radius
distance between nucleus and its valence shell
left to right: decreases
- pulls electrons closer to nucleus so atoms become smaller
- nuclear charge increases b/c atomic radius decreases
bottom to top: decreases
- less shells
27
# definition and pattern
atoms with low IE
atoms with high IE
noble gasses
ionization energy
energy require to remove an electron from a gaseous atom or ion
left to right: increase
- more protons
bottom to top: increase
-n bigger (bottom) means electrons are further from nucleus so protons can't hold onto the electrons as well
## Footnote
* atoms with low IE form cations
* atoms with high IE form anions
* noble gases have a very high IE so they're very stable and have a full shell
28
# definition and pattern
electron affinity
energy change when an electron is added to a gaseous atom
left to right: increase
- more protons so stronger attraction for electrons added
* bottom to top: increase
- easier to add electron at lower energy b/c less stable
## Footnote
* think of it as energy required to gain an electron
* opposite of ionization energy
* atoms with low EA tend to form cations
- low EA (low attraction for additional electrons)
* atoms with high EA tend to form anions
29
as nuclear charger increases...
ionization energy increases
30
as the distance between the electron and the nucleus increases...
ionization energy decreases
31
as shielding increases...
ionization energy decreases
## Footnote
outer electrons experiences less pull from the nucleus
32
rule for removing electrons
remove from the highest energy first
33
ionic radius
radius of a cation or ion
## Footnote
left to right: decreases
- more protons pulling electrons closer
bottom to top: decreases
- less shells
34
cations
| smaller or bigger than their atoms
positively charged ions from electrons getting removed
## Footnote
smaller than their atoms b/c the protons have to pull less electrons closer to nucleus → higher nuclear charge → decreasing ionic radius
35
anions
larger than their atoms b/c more electron-electron repulsion and more electrons to pull
36
isoelectronic ions
same # electrons
* have different size radii b/c ionic radius depends on protons too
37
compound
sunstance composted of 2+ elements in a specific ratio and held together by chemical bonds
38
ionic compound
pure substance formed from a metal and a nonmetal and has an overall neutral charge
39
covalent bond
sharing of outermost electrons
40
law of conservation of mass
mass is constant during ordinary chemical reactions
41
law of definite proportions
ex. water is always H2O
42
law of multiple proportions
some elements combine w/ each other in different whole number proportions
## Footnote
ex. CO2 vs CO
43
# definition and formula
molecular weight
sum of atomic masses
44
percent composition
n x molar mass of the element/molar mass of compound x 100
45
empirical formula
the lowest whole number ratio of the atoms present in one molecule of the compound
## Footnote
steps:
1. assume 100 g and find moles for each element
2. write it out
3. divide by smallest mole
46
molecular formula
number of each type of atom present in one molecule in the compound
## Footnote
1. find n using n=molar mass/mass of EF
2. multiply empirical formula by n
47
structural formula
arrangment of the atoms in the molecule
## Footnote
based on molecular formula
48
which elements can form double bonds
C, O, N, and sometimes S
49
bonding capacity
number of bonds an element can form (single elctrons)
50
# definition and pattern
electronegativity
ability to attract electrons
## Footnote
* left to right: increases
- more protons
* bottom to top: increases
- atomic radius is smaller → protons have a better pull on electrons
- less shells → outer electrons have less inner electrons to shield them from protons
51
Is S or N more electronegative and why?
N b/c its smaller which outweighs the fact that S is more to the right
52
Formal charge
number of valence electrons - (# of nonbonding electrons+#of bonds)
## Footnote
* favorable if all F.C.= 0 or most
* negative has to be on most electronegative atom
53
exceptions to the octet rule
1. electron deficient
- Be, H, He, and B are more stable incomplete
- F is better w/ single bonds
2. free radicals
3. expanded octet
- atoms in periods 3-7 are capable of exceeding 8 electrons and have empty d-orbitals that can hold extra V.E
54
# definition and formula
Lattice energy
energy change when gaseous ions react to form a solid ionic compound
| compare charges
if charges are the same, look at atomic radius
## Footnote
L.E = kQ1Q2/R
Q↑ LE↑
R↑ LE ↓
55
bond energy
energy needed to break the attraction between nuclei and shared electrons
56
bond length
shorter=stronger
more=stronger
57
electron delocalization
resonance
58
isomers
have same M.F. but different bonding between atoms
| different molecules
59
resonance structures
differ in position of electrons
| same molecule
60
VSEPR
| valence shell electron pair repulsion
all electrons around central atom arrange themselves to be far away from each other as possible to minimize electronic repulsions
* predicts molecular shape
61
conditions for a polar molecule
1. at least 1 polar bond (can have polar bonds and be nonpolar)
2. have an asymmetrical shape
## Footnote
if dipoles don't cancel
62
polar molecules
* 2 polar bonds <180 degrees
* 3 polar bonds
* 3 polar bonds in trigonal pyramidal arrangment (109.5)
63
nonpolar molecules
* opposite directions
* all 120 degrees
* 4 identical polar bonds in tetrahedral arrangment (109.5)
64
hybridization
combining atomic orbitals to make hybrid orbitals
65
valence bond theory
a covalent bond forms when orbitals of 2 atoms overlap and a pair of electrons ocuupy the overlap region
## Footnote
lowers potential energy for the system
66
# valence bond theory continued
a bond will form if...
1. orbitals of 2 atoms overlap
2. total # of electrons is no more than 2
3. greater the overlap = stronger the bond
4. electrons are attracted to both nuclei, pulling atoms together
outer e- drawn to both nuclei
| inner e- drawn to nucleus of its atom
67
hybridization steps
1. draw ground state orbital diagram for **central atom**'s valence electrons
2. maximize number of unpaired electrons by promotion
3. combine necessary number of atomic orbitals to generate required number of hybrid orbitals
4. place electrons in hybrid orbitals
## Footnote
of electron domains dictates how many unhybridized orbitals
68
sigma bonds
* single bonds
* covalent bond where electrons are shared between atoms and the overlap region is **between** 2 nuclei
- b/c VE need to be close to nuclei of both atoms
* electron density is concentrated along internuclear axis
69
pi bonds
* formed from 2 p-orbitals overlapping
* each pi bond has 2 lobes/halves
* each electron could be both above and below
* no density along axis
70
rotation
| sigma and pi bonds
* sigma bonds have free rotation around bonx axis
- if multiple sigma bonds, each part can rotate in diff. directions
* pi bonds rotation is restricted
71
intramolecular forces
hold atoms together within a molecule
72
Are intramolecular or intermolecular forces weaker?
intermolecular forces
73
types intermolecular forces
1. dipole-dipole
2. H-bonding
3. london dispersion forces
4. ion interactions
74
dipole-dipole
* between polar molecules
* larger the dipole → larger the attractive force → higher the boiling point (takes more energy to break bonds to become gas)
75
Hydrogen bonding
| Hydrogen bonding donor and accepter
## Footnote
talk about boiling point
type of dipole dipole interaction when H-bonds w/ flourine, oxygen, or nitrogen
donor: donates H+ and is the entire molecule the hydrogen is apart of
acceptor: accepts H+
| more H-bonds → higher boiling point
## Footnote
F, O, N have to have lone pairs
H has to be attached to a more EN atom (also F, O, N)
76
Dispersion forces
* occurs in all molecules (only one that occurs in 2 polar molecules)
* temporary dipoles
* includes london forces and induced dipoles
77
induced dipoles
1. 2 nonpolar (london)
2. 1 polar and 1 nonpolar
3. 2 polar
4. ion and nonpolar
## Footnote
as negative end of polar molecule approaches nonpolar molecule, the electrons in nonpolar move away to reduce repulsion
78
ion interactions
between ions and polar or ions and nonpolar
coloumbic so short distance → stronger dipole
79
bond strength order
80
which groups attract water and which ones don't
polar and ionic do
nonpolar doesn't
81
ionic compounds
| intermolecular force
solvent cage
dissociate into ions when they dissolve
ion dipole
solvent cage: group of solvent molecules that surround a dissolved ion in a solution
82
miscible
"like dissolves like"
2 liquids are miscible if they're soluble in all properties
## Footnote
* polar and nonpolar are imiscible
