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MCAT chem > atomic trends and bonding/ IMF > Flashcards

atomic trends and bonding/ IMF Flashcards

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1
Q

what is all matter composed of?

A

atoms

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2
Q

what is a molecule?

A

when two or more atoms are combined

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3
Q

what is a compound

A

when two or more molecules are combined

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4
Q

proton charge and weight

A

at center of atom, weigh 1 amu and have 1+ charge
1.6E-19 eV

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5
Q

electron weight and charge

A

in electron clouds, -1 or -1.6E-19

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6
Q

neutron weight

A

1 amu (same as proton_

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7
Q

atomic number (z)

A

identifies element, number of protons

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8
Q

mass number

A

P + N

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9
Q

isotopes

A

same atomic number but differnt number of neutrons

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10
Q

atomic mass

A

N+P

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11
Q

automic weight

A

average of all isotopic automic masses

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12
Q

polyatomic ions

A

contain multiple atoms with different ionic forms

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13
Q

oxyanions

A

polyatomics that contain oxygen and have suffix ite or ate

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14
Q

bohr model

A

outdated but explains how electrons exsist in clouds outside the nucleus in shells

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15
Q

ground state

A

lowest energy level of an electron n=1,2,3ā€¦ and can jump to higher energy levels

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16
Q

photons

A

discrete amounts of energy that are emitted as electrons jump to excited states

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17
Q

energy of electromagnetic ratiation equation

A

E= hf where h is planks constant and f is the frequency of light (which is the speed of light divided by wave length)
E= h*c/ wavelength

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18
Q

rydberg formula
used to determine the energy held by an electron in a certain energy level

A

change in Energy = R (1/nf^2 - 1/ni^2)
r is constant

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19
Q

Heisenberg principle of uncertainty

A

cannot know the momentum and exact position of an electron at the same time

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20
Q

principle quantum number

A

electrons exist in orbitals
denoted as ā€œnā€ and can be 1 or greater, higher n = higher energy and farther from the nucleus
- associated with rows of the periodic table

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21
Q

orbital

A

area of space that electrons are likely to be in, can only hold two electrons at a time

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22
Q

pauli exclusion principle

A

no two electrons in an atom can have the exact same values for all 4 quantum numbers

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23
Q

azimuthal/angular momentum quantum number

A

describes the subshell
s,p,d,f

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24
Q

magnetic quantum number

A

describes the spacial orientation of the orbital within its subshell ranging from + and - L. L=0 for s, L=1 for p, L=2 for d L= 3

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25
spin quantum number
describes the spin orientation of the electron either +1/2 or -1/2. all electrons are paired with opposite spin
26
electron configuration
configuration of electrons in subshells , follow aufbau principle and lower energy levels fill first
27
hunds rule
electrons prefer to be in orbitals by themselves and will fill each orbital with 1 electron before sharing
28
where is the s block
the first 2 groups
29
where is the p block
groups 13-18
30
where is the d block
transition metals
31
where is the f block
lanthanide and actinide series at the bottom of the periodic table
32
valence electrons
outermost electrons that are involved in chemical bonds
33
which is a lower energy? 3d or 4s
4s so it fills before 3d (according to aufbau principle)
34
true or false: half filled and full filled orbitals are more stable than ones with 3/4/ or 9 electrons
true! the p and d will steal from s orbital
35
alkali metals
group 1: high conductivity, luster, duct-ability (drawn into thin wire) , malleability, reactivity, reactive
36
alkaline earth metal
group 2: harder and less reactive than alkali metals, good conductors
37
transition metals
groups 3-12. luster, hard, conduct electricity
38
nonmetals
not shiny, poor electricity conductors
39
chalcogens
group 16, feature O and S
40
halogens
group 17, highly reactive
41
noble gases
group 18, very stable. inert gases. low boiling points
42
effective nuclear charge
the attraction between the positively charged center of an atom and the electrons surrounding it. decreases down a group and increases to the right. increases with increased number of protons
43
atomic radius trend
increases down a group and to the left
44
ionic radius
cations are smaller than anions
45
ionization energy
the amount of energy required for an atom to give up an electron. increases across a period and up a group
46
electron affinity
the amount of energy released when an electron is added to an atom. increases up a group and to the right.
47
electronegativity
tendency of an atom to attract electrons, directly in correlation with effective nuclear charge
48
how many molecules are in one mole?
6.02E23 molecules
49
octet rule
each atom will act in a way that allows it to obtain 8 electrons in its valence shell for stability.
50
is bond formation typically endothermic or exothermic?
it is stable and releases energy so it is exothermic
51
incomplete octet
special case for hydrogen where it is a filled valence shell with a single bond H-H others: Boron - 6 he- 2 li-2 be - 4
52
expanded octet
atom has more than 8 valence electrons. in case of PF5 there are 10
53
what is unique about octets in the third period?
they readily form expanded octets
54
intramolecular forces (definition and types)
forces that work inside a molecule to keep it together ionic and covalent bonds
55
ionic bonds
form between a metal and a nonmetal, the two species have a large difference in ELECTRONEGATIVITY. form cations and anions
56
pauling scale
assigns electronegativity values. fluorine is the most eneg @ 4.0 and the alkali metals have an eneg of around 1.0 ** ionic bonds form when the difference in electronegativity between two atoms is 1.7 or greater
57
covalent bonding
sharing of electrons among nonmetals - can be equal or unequal sharing - the difference in electronegativity between the two atoms is small
58
nonpolar covalent bonds
when electronegativity between atoms sharing electrons are similar or the same and they share electrons equally - <.5
59
polar covalent bonds
occur when difference in electronegativity is between .5 and 1.7 and electrons are shared unequally, with the more electronegative atom recieving more electrons and causing dipole moments
60
dipole moments
unequally sharing of electrons causes partial positive charges on less electronegative atoms pointing toward the more electronegative atom with a partial negative
61
coordinate covalent bond
when shared electrons are derived from a lone pair of one atom
62
bond order
number of bonds between two atoms single, double, triple
63
sigma bond
forms due to end to end overlap of two atomic orbitals a single bond forms 1 sigma bond
64
pi bond
forms when atomic orbitals point above and below the atoms making them parallel. a double bond contains a pi and a sigma
65
bond length
inversely related to bond order. single bonds are the longest and triple bonds are the shortest
66
bond energy
energy required to break a covalent bond. the higher the bond energy, the stronger the bond - increases with bond order
67
metallic bonding
an intramolecular force less common than ionic and covalent. occurs when metal atoms are joined together and electrons become delocalized in a sea of electrons that freely move through a solid and allows for electrical conduction and heat
68
intermolecular forces
attractive forces between molecules that are notably WEAKER than intramolecular forces and can be broken with heat . bonds driven by charge types: london, dipole, hydrogen, ion-dipole - stronger polarity = stronger intermolecular forces - determine behavior and melt/boiling points
69
london-dispersion forces
weakest intermolecular force caused by random dipole moments, fluctuations in polar and nonpolar bonds - larger compounds are more likely to experience london dispersion forces
70
dipole-dipole interactions
intermolecular force that is stronger than london dispersion. dipoles are stable (polar bonds only)
71
hydrogen bonding
between a hydrogen bound to F, O, N and a lone pair on another F, O, or N
72
ion-diple forces
strongest intermolecular force. ions are most polar and occur when two molecules have charge ex: peptide hormone signalling with a basic aminoacid
73
lewis structure
establish relationships between atoms and their valence electrons - valence electrons = group number (except transition metals) - atom with lowest electronegativity goes in the middle
74
resonance structures
compounds with lone pairs that can move bonds to different positions - can only shift lone pairs and double bonds, never single bonds - delocalization of electrons increases stability of structure
75
formal charge
disparity between the number of valence electrons an atom has and the electrons it actually has. formal charge = VE - bonds - lone pair electrons
76
orbital hybridization
mixing atomic orbitals to form hybrid orbitals
77
what does a triple bond consist of
2 pi and 1 sigma bond
78
two regions of hybridization (triple bond) yields what?
sp
79
three regions of hybridization (double bond) yields what?
sp2
80
4 regions of hybridization (single bond) yields what?
sp3
81
VSPER theory
uses lewis structures to determine shapes of molecules due to electron repulsion - use number of lone pairs and bonded atoms to predict bond angle and shape
82
electronic geometry
takes into account lone pairs and bonded atoms to predict shape
83
molecular geometry
only takes into account boned atoms when predicting shape
84
tetrahedral molecular shape
4 electron rich regions and 4 bonded atoms - CH4 - bond angle is 109.5
85
trigonal pyrimidal molecular shape
4 regions of electron rich. 3 bonded atoms and 1 lone pair - bond angle is 107 nh3
86
bent molecular shape
2 bonds and 2 lone pairs, bond angle is 104 ex: h2o
87
linear molecular shape
2 bonds no lone pairs, bond angle is 180
88
trigonal planar molecular shape
3 bonds no lone pairs, bond angle 120 ex bf3

MCAT chem (11 decks)

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