Chapter 1 General Chemistry Review Flashcards
1
Q
define organic chemistry
A
chemistry of compounds containing carbon
2
Q
define inorganic chemistry
A
chemistry of compounds containing metals
3
Q
define organometallic chemistry
A
chemistry of compounds containing carbon and metals
4
Q
organic chemistry studies the distribution and motion of what?
A
electrons
5
Q
how do electrons move?
A
from negative regions (areas of electron density) to positive regions (areas of electron deficiency)
6
Q
what can block electron flow, and what does this mean?
A
electron flow can be blocked by large groups, so shape and size of a molecule is important
7
Q
what happens when electrons move?
A
a chemical reaction
8
Q
what do reaction mechanisms represent?
A
how electrons are thought to flow in a chemical reaction
9
Q
describe an important aspect (+ and -) of a reaction mechanism
A
negative charges (high electron density) attack positive charges (low electron density)
10
Q
what 3 things must be considered to understand the structure of organic compounds?
A
- what atoms does the molecule contain?
- how are the atoms connected?
- what types of bonds connect the atoms?
11
Q
by what two things are compounds defined?
A
- type and number of atoms (molecular formula)
2. how those atoms are connected to each other (connectivity/constitution)
12
Q
define constitutional isomers
A
possess the same type and number (molcular formular) of atoms but are connected DIFFERENTLY
13
Q
give the 2 steps for determining a constitutional isomer
A
- check MF, if same keep going, if not same, not constitutional isomers
- check connectivity, must be different
14
Q
from what three things do the properties of an organic compound arise?
A
- the atoms they contain
- how they are connected
- the type of bonds contained
15
Q
what does different connectivity give constitutional isomers?
A
different chemical and physical properties
16
Q
what holds atoms together in molecules?
A
chemical bonds
17
Q
what do an atom’s properties determine?
A
its bonding
18
Q
elements are defined by the number of?
A
protons
19
Q
what is the atomic number (Z)?
A
number of protons
20
Q
what is mass number (A)?
A
number of protons plus number of neutrons
21
Q
describe protons and neutrons in a neutral atom
A
same amount of each
22
Q
for an uncharged atom, what does the number of electrons equal?
A
the number of protons
23
Q
where do electrons reside in atoms?
A
in atomic orbitals
24
Q
how many atomic orbitals are there? list them
A
4: s, p, d, f
25
what two things are atomic orbitals defined by?
quantum mechanics and Schrodinger wave equation
26
what 2 properties do electrons exhibit?
particle and wavelike
27
define atomic orbital
region of space that can be occupied by electron density
28
what is an occupied atomic orbital compared to?
a "cloud" of electron density
29
what are the shapes of atomic orbitals?
s, p, d, f
30
what does the shape of an atomic orbital represent?
the region of space that contains 90-95% of the electron density
31
what are phases of waves?
regions of positive and negative amplitude
32
can a wave or electron ever be in the middle of amplitude?
no, only positive or negative
33
what is the region in the middle of the wave (and amplitude) called?
node
34
describe the regions of an electron orbital in its wavelike property
have postive and negative regions (phases) and nodes
35
describe the node of an atomic orbital?
where there is zero probability of finding an electron, or zero electron density
36
what do electron configurations list?
the orbitals containing electrons and the number of electrons in each orbital
37
define core electrons
all electrons in completely filled inner energy shells (lower energy)
38
define valence electrons
electrons in highest energy outer shell (higher energy = REACTIVE)
39
when are atoms most stable?
when they have a full valence shell
40
what is a full valence shell also referred to as?
a noble gas configuration
41
what are noble gases stable (unreactive)?
they have a full valence shell
42
how do second row atoms (C, N, O, F) achieve a full valence shell?
by obtaining an OCTET of electrons by forming bonds
43
what is an atomic energy diagram?
when an electron configuration is drawn to show relative energies of the filled or empty orbitals
44
what are individual atomic orbitals within the same orbital subshell called? what does that mean?
degenerate, same energy
45
list the 7 steps for drawing an atomic energy diagram
1. draw a line to represent orbitals
2. lower energy orbitals are drawn at the bottom, higher energy orbitals are drawn at the top
3. use periodic tabe to determine which orbitals to draw
4. orbitals in the same subshell (3 p or 5 d) are degenerate
5. Aufbau principle
6. Pauli exclusion principle
7. Hund's rule
46
define the Aufbau principle
fill lowest energy orbitals first, then proceed to higher energy orbitals
47
define Pauli's exclusion principle
each orbital can contain 2 electrons which must have opposite spin (spin up or down)
48
define Hund's rule
for degenerate orbitals one electron is placed in each orbital before electrons are paired (minimizes repulsion of electrons giving lower energy arrangement)
49
define a chemical bond
shared electron pair between two atoms
50
what does sharing a pair of electrons allow each atom in the bond to do?
gain an additional electron
51
why do atoms form chemical bonds?
to achieve a full valence shell
52
what does achieving a full valence shell do to an atom?
stabilizes it
53
how does hydrogen acieve a full valence shell and why?
with only 2 electrons with only 2 electrons since it has only the 1s orbital
54
define covalent bond
each atom donates 1 electron and the 2 atoms share the electron pair, each atom obtains a full valence shell!
55
define ionic bond
attraction between 2 oppositely charged ions. one atom possesses both electrons in the pair
56
considering Hydrogens electron configuration (1s1) how many bonds is it expected to form?
one bond
57
considering heliums electron configuration (1s2) how many bonds is it expected to form?
none
58
with carbons eletron configuraiton (1s2,2s2,2p2) how many bonds is it expected to form?
4
59
with nitrogens electron configuration (1s2,2s2,2p3) how many bonds is it expected to form
3
60
with oxygens (1s2,2s2,2p4) electron configuration how many bonds is it expected to form
2
61
with fluorines electron configuration (1s2,2s2,2p5) how many bonds is it expected to form
1
62
when counting the total valence electrons on an atom in a molecule how much do bonds and lone pairs count for?
bonds = 2 electrons
| lone pairs = 2 electrons
63
how is hydrogen an exception to the valence shell octet?
it will only form one bond to achieve the noble gas configuration of helium, filling 1s orbital only
64
how is boron an exception to the valence shell octet?
boron only has 3 valence shell electrons and therefore will only form 3 bonds. this gives boron only 6 valence electrons, and makes compounds containing boron especially reactive because of the incomplete octet
65
define expanded valences
elements in the 3rd period (row) and below (especially P and S) often possess 10 or 12 valence electrons since they have a 3rd orbital which is capable of holding electrons
66
describe phosporous (P) as it acts as an exception to the valence shell octet
typically forms 3-5 bonds
67
describe sulfur (S) as it acts as an exception to the valence shell octet
typically forms 2, 4, or 6 bonds
68
why and how do atoms form bonds?
to gain stability through filling their valence shell
69
how many bonds will atoms want to form?
the number required to fill their valence shell (with exceptions)
70
how do lewis structures represent molecules?
draw the bonds in molecules by only considering valence electrons (outermost electron shell)
71
what does a line between 2 atoms represent in a lewis structure?
a bond (2 electrons)
72
how are nonbinding electrons shown in a lewis structure?
as 2 dots, or a lone pair
73
give the 5 steps for drawing a lewis structure
1. draw lewis structures of all individual atoms
a. place atomsforming more than one bond in the center
b. place atoms forming only one bond (H, halogens) on exterior
2. connect atoms forming more than 1 bond
3. connect monovalent atoms (H and halogens)
4. check to see if all atoms have a full valence shell
a. bonds = 2 valence electrons
b. lone pairs = 2 valence electrons
c. H = 2 electrons only
5. if atoms lack full valence shell combine additional unpaired electrons to form double and triple bonds to achieve full valence shell (usually octet)
74
give the 2 cases in which an atom would bear a formal charge
1. atom has more electrons than the number of protons in the atom= negative charge
2. atom has less electrons than the number of protons in the atom = positive charge
75
when an atom gains valence electrons, what is its new charge and what is it called?
negative charge, anion
76
when an atom loses valence electrons, what is its new charge and what is it called?
positive, cation
77
what are the 2 questions you should ask when assigning formal charge? how do you answer them?
1. how many valence electrons should the atom have?
a: check group (column) in periodic table
2. how many formal charge electrons reside on the atom in the lewis structure?
a: lone pairs = 2 electrons, bonds = 1 electron
78
give the formula for calculating formal charge
formal charge = number of valence electrons from periodic table - (number of bonds + number of lone pair electrons)
79
when are atoms neutral?
when the number of bonds on the atom = the number of unpaired valence electrons
80
what does removing a bone and adding a lone pair do to the formal charge of an atom and why?
introduces a negative charge since this causes the atom to gain an electron
81
what does adding a bond and removing a lone pair do to the formal charge of an atom and why?
introduces a positive charge since the atom "loses" an electron by sharing the electrons in the bond
82
when drawing lewis structures of molecules with a net formal charge what should you add for negatively charged molecules?
add an extra electron
83
when drawing lewis structures of molecules with a net formal charge what should you remove for a postively charge molecule?
remove an electron
84
what are the 2 main bonding theories used in ochem to explain molecular structure and reactivity?
1. valence bone theory: VSEPR and hybridization
| 2. molecular orbital theory
85
what does VSEPR stand for?
valence shell electron repulsion theory
86
in terms of valence bond theory, what is a covalent bond?
overlap of 2 atomic orbitals (atomic orbitals have wavelike phases)
87
what does overlap of 2 waves with the same phase produce?
contructive interference, resulting in a wave with alrger amplitude, A BONDING INTERACTION
88
what does overlap of waves with opposite phases create?
destructive interference, cancelling each other out and producing a node
89
what is a node?
a region of no electron density, anti-bond interaction
90
what are the 2 types of bonds?
1. sigma
| 2. pi
91
what is a sigma bond?
head to head overlap of two orbitals
92
where does electron density exist in a sigma bond?
between bonding nuclei along internuclear axis
93
what kind of symmetry exists around sigma bonds and what does this mean?
circular symmetry around bond axis = no nodes
94
what is a pi bond?
side-to-side overlap of two P orbitals (P for pi bond)
95
wheredoes electron density exist in a pi bond?
above and below bonding nuclei
96
where does the node exist in a pi bond?
along the internuclear axis
97
what happens with VSEPR?
negatively charged electrons in bonds and lone pairs (nonbonding electrons) arrange as far apart as possible to minimize repulsion
98
what is steric number?
the total number of sigma bonds and lone pairs (single, double, and triple bonds count as 1 for steric number)
99
what does steric number determine?
the arrangement of electron pairs (lone pairs and sigma bonds)
100
what does the number of atoms determine in VSEPR?
geometry
101
does molecule geometry include electrons?
NO only atoms for geometry
102
what does an atom need in order to have geometry?
to be connected to at least TWO other atoms
103
do atomic orbitals minimize electron repulsion in molecules?
no, only minimize electron repulsion in atoms
104
what do atomic orbitals do to minimize electron repulsion in MOLECULES?
combine to form new shapes (HYBRIDIZE)
105
what do you need to form new identical hybridized bonds?
degenerate orbitals
106
what does an unsymmetrical hybridized sp3 orbital look like?
larger front lobe, so will often omit smaller lobe in drawings
107
is a pi bond or a sigma bond stronger?
sigma bond stronger
108
which will react more, a pi or sigma bond, and why?
pi bond will react more because it is not as strong, can be broken more easily
109
order single, double, and triple bonds according to bond strength
triple bond > double bond > single bond
110
order single, double, and triple bonds according to length
single bond > double bond > triple bond
111
is a double bond twice as strong as a single bond? why or why not?
no, because a pi bond is not as strong as a sigma bond due to the node along the internuclear axis
112
what is electronegativity?
the measure of an atom's ability to attract electrons
113
what does a higher electronegativity value mean for an atom?
greater electron attraction power (pulls electrons to itself)
114
what is the most electronegative element?
fluorine (F)
115
what 2 things affect electronegativity?
1. positive character of atom's nucleus
| 2. shelding of nucleus by electrons
116
why does the positive character of an atom's nucleus effect electronegativity?
additional protons (positive charge) increases an atom's electronegativity
117
how does electronegitivity trend across the periodic table?
increases from left to right and down to up across the periodic table
118
how does shielding of nucleus by electrons effect electronegativity?
additional shells of electrons shield the positive character of the nucleus and decrease electronegativity of the atom
119
what is induction?
withdrawal of electron density through a bond bewteen 2 atoms of DIFFERENT electronegativity
120
what does induction create? 2 names
a polar bond or "bond dipole moment"
121
what is a dipole moment?
a separation of charge over some distance
122
what does a separation of charge over distance result in?
a polar bond!
123
give the 3 bond categories
1. covalent
2. polar covalent
3. ionic
124
define a covalent bond (EN# range and definition)
EN < 0.5, electrons shared EVENLY between two atoms
125
what is an example of a covalent bond?
C-H
126
is a C-H bond polar?
NOOOOOO
127
define a polar covalent bond (EN # range and definition)
EN 0.5-1.7, increased electron density on more electronegative atom
128
give three examples of polar covalent bonds
C-O, C-N, and C-X (where X = a halogen)
129
define an ionic bond (EN# range and definition_
EN > 1.7, more electronegative atom possesses ALL electron density
130
give an example of an ionic bond
Na-OH
131
describe the charges when you see an atom bound to a metal and why
the atom bound to the metal has a negative charge! metals are very electroPOSITIVE, so they will give up their electrons
132
what is a molecular dipole moment?
the vector (direction and magnitude) sum of the individual bond dipole moments
133
what MUST you consier when calculating molecular dipole moment?
the geometry of the molecule
134
what are dipole moments associated with?
electron lone pairs
135
where does the dipolemoment point when an electron lone pair is involved?
toward the lone pair
136
what are the physical properties of a compound determined by?
its intermolecular forces that are electrostatic (attraction of opposite charges) that result from molecular dipole moments and individual dipole moments
137
give the 3 intermolecular forces in order of decreasing strength
1. hydrogen bonding
2. dipole-dipole interactions
3. london dispersion forces
138
what do stronger intermolecular forces (stronger attraction between molecules) result in?
higher boiling points and higher melting points
139
describe hydrogen bonding (4)
1. interaction of partial positive hydrogen with LONE PAIR of an electronegative atom (usually N or O)
2. partial hydrogen occurs when hydrogen is bound to an electronegative atom
a. hydrogen takes on positive character due to bond with electronegative atom
3. not a real bond
4. strong interaction due to small size of hydrogen allowing shorter intermolecular distance between H-bonding partners (4-50 kJ/mol "bond" strength)
140
describe dipole-dipole interactions (3)
1. interaction of partial positive region of molecule woth partial negative region of molecule
2. molecule must possess a net dipole to exhibit dipole-dipole interactions
3. weaker than H-bonding
141
describe London Dispersion forces
1. even if a molecule does not possess a molecular dipole moment, its electrons are constantly in motion, giving rise to fleeting dipole moments in neighboring molecules
2. momentary interaction of transient or fleeting partial positive region of molecule with transient partial negative region of another molecule are know as london dipersion forces
3. weakest intermolecular force
4. larger compounds have more surface area and experience these forces to a greater degree
5. branched compounds have a smaller surface area, therefore experience less london dispersion forces
