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CHM 129 > Exam 1 > Flashcards

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

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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

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3
Q

average atomic mass

A

weighted average of all the naturally occuring isotopes of the element

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4
Q

atomic number

A

protons
number on periodic table above the element symbol

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5
Q

mass number

A

not on periodic table
protons + neutrons

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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

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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
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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
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9
Q

orbital

A

locations in space around nucleus at which the probably of finding the electron has higher values

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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

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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

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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
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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

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14
Q

electron spin (ms)

A

ms = +1/2, -1/2

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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)

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16
Q

number of electrons in a shell formula

A

2(n)^2

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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

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18
Q

shielding

A

inner electrons block outer electrons from protons

  • electrons in the same energy level shield each other
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19
Q

Aufbau Process

A
  1. electrons occupy the lowest energy possible
  2. Hund’s Rule: every orbital is singly occupied w/ one electron before any are doubly occupied
  3. Pauli exclusion principle: no 2 electrons in an atom have the same 4 quantum numbers
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20
Q

Aufbau process exceptions

A

transition metals
lanthanides
actinides

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21
Q

how do you know which could represent the electron in the highest energy shell

A

the coefficient

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22
Q

inner core electrons

A

those an atom has in common w/ the previous noble gas

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23
Q

outer electrons

A

those in the highest energy level (highest n-value)

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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)
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25
Q

definition and pattern

nuclear charge (Zeff)

A

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

  • 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.
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26
Q

definition and pattern

atomic radius

A

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

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27
Q

definition and pattern

atoms with low IE
atoms with high IE
noble gasses

ionization energy

A

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

  • 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
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28
Q

definition and pattern

electron affinity

A

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

  • 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
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29
Q

as nuclear charger increases…

A

ionization energy increases

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30
Q

as the distance between the electron and the nucleus increases…

A

ionization energy decreases

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31
Q

as shielding increases…

A

ionization energy decreases

outer electrons experiences less pull from the nucleus

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32
Q

rule for removing electrons

A

remove from the highest energy first

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33
Q

ionic radius

A

radius of a cation or ion

left to right: decreases
- more protons pulling electrons closer
bottom to top: decreases
- less shells

34
Q

cations

smaller or bigger than their atoms

A

positively charged ions from electrons getting removed

smaller than their atoms b/c the protons have to pull less electrons closer to nucleus → higher nuclear charge → decreasing ionic radius

35
Q

anions

A

larger than their atoms b/c more electron-electron repulsion and more electrons to pull

36
Q

isoelectronic ions

A

same # electrons
* have different size radii b/c ionic radius depends on protons too

37
Q

compound

A

sunstance composted of 2+ elements in a specific ratio and held together by chemical bonds

38
Q

ionic compound

A

pure substance formed from a metal and a nonmetal and has an overall neutral charge

39
Q

covalent bond

A

sharing of outermost electrons

40
Q

law of conservation of mass

A

mass is constant during ordinary chemical reactions

41
Q

law of definite proportions

A

ex. water is always H2O

42
Q

law of multiple proportions

A

some elements combine w/ each other in different whole number proportions

ex. CO2 vs CO

43
Q

definition and formula

molecular weight

A

sum of atomic masses

44
Q

percent composition

A

n x molar mass of the element/molar mass of compound x 100

45
Q

empirical formula

A

the lowest whole number ratio of the atoms present in one molecule of the compound

steps:
1. assume 100 g and find moles for each element
2. write it out
3. divide by smallest mole

46
Q

molecular formula

A

number of each type of atom present in one molecule in the compound

  1. find n using n=molar mass/mass of EF
  2. multiply empirical formula by n
47
Q

structural formula

A

arrangment of the atoms in the molecule

based on molecular formula

48
Q

which elements can form double bonds

A

C, O, N, and sometimes S

49
Q

bonding capacity

A

number of bonds an element can form (single elctrons)

50
Q

definition and pattern

electronegativity

A

ability to attract electrons

  • 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
Q

Is S or N more electronegative and why?

A

N b/c its smaller which outweighs the fact that S is more to the right

52
Q

Formal charge

A

number of valence electrons - (# of nonbonding electrons+#of bonds)

  • favorable if all F.C.= 0 or most
  • negative has to be on most electronegative atom
53
Q

exceptions to the octet rule

A
  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
Q

definition and formula

Lattice energy

A

energy change when gaseous ions react to form a solid ionic compound

compare charges
if charges are the same, look at atomic radius

L.E = kQ1Q2/R
Q↑ LE↑
R↑ LE ↓

55
Q

bond energy

A

energy needed to break the attraction between nuclei and shared electrons

56
Q

bond length

A

shorter=stronger
more=stronger

57
Q

electron delocalization

A

resonance

58
Q

isomers

A

have same M.F. but different bonding between atoms

different molecules

59
Q

resonance structures

A

differ in position of electrons

same molecule

60
Q

VSEPR

valence shell electron pair repulsion

A

all electrons around central atom arrange themselves to be far away from each other as possible to minimize electronic repulsions
* predicts molecular shape

61
Q

conditions for a polar molecule

A
  1. at least 1 polar bond (can have polar bonds and be nonpolar)
  2. have an asymmetrical shape

if dipoles don’t cancel

62
Q

polar molecules

A
  • 2 polar bonds <180 degrees
  • 3 polar bonds
  • 3 polar bonds in trigonal pyramidal arrangment (109.5)
63
Q

nonpolar molecules

A
  • opposite directions
  • all 120 degrees
  • 4 identical polar bonds in tetrahedral arrangment (109.5)
64
Q

hybridization

A

combining atomic orbitals to make hybrid orbitals

65
Q

valence bond theory

A

a covalent bond forms when orbitals of 2 atoms overlap and a pair of electrons ocuupy the overlap region

lowers potential energy for the system

66
Q

valence bond theory continued

a bond will form if…

A
  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
Q

hybridization steps

A
  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

of electron domains dictates how many unhybridized orbitals

68
Q

sigma bonds

A
  • 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
Q

pi bonds

A
  • 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
Q

rotation

sigma and pi bonds

A
  • 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
Q

intramolecular forces

A

hold atoms together within a molecule

72
Q

Are intramolecular or intermolecular forces weaker?

A

intermolecular forces

73
Q

types intermolecular forces

A
  1. dipole-dipole
  2. H-bonding
  3. london dispersion forces
  4. ion interactions
74
Q

dipole-dipole

A
  • between polar molecules
  • larger the dipole → larger the attractive force → higher the boiling point (takes more energy to break bonds to become gas)
75
Q

Hydrogen bonding

Hydrogen bonding donor and accepter

talk about boiling point

A

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

F, O, N have to have lone pairs
H has to be attached to a more EN atom (also F, O, N)

76
Q

Dispersion forces

A
  • occurs in all molecules (only one that occurs in 2 polar molecules)
  • temporary dipoles
  • includes london forces and induced dipoles
77
Q

induced dipoles

A
  1. 2 nonpolar (london)
  2. 1 polar and 1 nonpolar
  3. 2 polar
  4. ion and nonpolar

as negative end of polar molecule approaches nonpolar molecule, the electrons in nonpolar move away to reduce repulsion

78
Q

ion interactions

A

between ions and polar or ions and nonpolar
coloumbic so short distance → stronger dipole

79
Q

bond strength order

A
80
Q

which groups attract water and which ones don’t

A

polar and ionic do
nonpolar doesn’t

81
Q

ionic compounds

intermolecular force
solvent cage

A

dissociate into ions when they dissolve
ion dipole
solvent cage: group of solvent molecules that surround a dissolved ion in a solution

82
Q

miscible

A

“like dissolves like”
2 liquids are miscible if they’re soluble in all properties

  • polar and nonpolar are imiscible

CHM 129 (4 decks)

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