Fundamentals of Chemistry Flashcards
1
Q
what is the charge and mass of a proton
A
positive, and 1 amu
2
Q
what is the charge and mass of a neutron
A
no charge, 1 amu
3
Q
what is the mass and charge of an electron
A
negative, 0.0005458 amu
4
Q
define ion
A
the number of protons does not equal the number of electrons
5
Q
define isotope
A
same number of protons and electrons but a different number of neutrons
6
Q
define pure substance
A
consists of atoms with the same number of protons, cannot be broken into simpler species
7
Q
define compound
A
consist of more than one type of element and are held together by chemical bonds
8
Q
define ionic compound
A
when a non-metal reacts with a metal and a transfer occurs. There is an electrostatic attraction between the opposite charges
9
Q
properties of ionic compounds
A
high melting point and can conduct electricity when liquid
10
Q
define molecular compound
A
covalent compounds where atoms share electrons, often non-metals
11
Q
properties of molecular compounds
A
low melting and boiling points
12
Q
define empirical formula
A
gives elements the smallest possible ratio
13
Q
define skeletal formula
A
each carbon is represented by a line
14
Q
define isomer
A
different molecules have the same formula but different connectivity
15
Q
define chemical reaction
A
any process that leads to a chemical transformation of one or more substance into another
16
Q
how many molecules does a mole contain
A
6.022x10^23
17
Q
define attractive force
A
one that pushes things towards zero (makes them closer)
18
Q
define repulsive force
A
pushes things away
19
Q
define potential well
A
a dip in energy vs distance graph
20
Q
define functional group
A
a part of a molecule which has distinctive chemical properties
21
Q
IR light will only be absorbed when shone at a sample ifβ¦
A
- the energy of the radiation corresponds to the energy required to vibrate the molecule
- the vibration leads to the dipole moment of the molecule changing
22
Q
what is the equation for electromagnetic radiation (speed of light)
A
wavelength x frequency
23
Q
what is the equation for energy of a photon
A
plancks constant x frequency of a photon
24
Q
define the Heisenberg Uncertainty Principle
A
we cannot simultaneously know the location and energy of an electron
25
define orbital
a region of the atom where an electron is most likely to reside
26
name the 4 different types of orbitals
different energies (n)
different shapes (l)
different orientations (ml)
different spins (ms)
27
what is n more commonly known as
the principle quantum number
28
what is l more commonly known as
the angular momentum quantum number (n-1)
29
what are orbitals with l=0 known as
s-orbitals
30
what are orbitals with l=1 known as
p-orbitals
31
what is an angular node
when the wave function changes sign, the point in the middle is 0, and this is an angular node, l
32
what are the values of ms
1/2 or -1/2
33
what did Pauli discover
no two electrons in the same atom can have exactly the same quantum numbers, they can share the same orbital provided ms is different, meaning there is a max of 2 electrons per orbital
34
list the orbitals from lowest to highest energy
s, p, d, f
35
describe metallic bonding
valence electrons are a long way from the nucleus and not held tightly, multiple atoms can overlap without nuclei repelling, does not require close overlap, electrons can delocalise across the system and carry charge
36
define power
the rate of change of energy
37
give the equation for power
E/t (energy/time)
38
define potential maximum
any change or movement will cause the energy to be lowered
39
define potential minimum
extra energy is needed to move it away from the minimum (most stable positions)
40
define Coulomb's Law
the force between two oppositely charged objects depends on the size of the charges and varies inversely with the square of the distance
41
if the overall charge is negative for Coulomb's equation, is the force attractive or repulsive
attractive
42
define binding energy
the energy given out by forming a bond
43
define crystalline solid
ordered at the microscopic scale, with the component atoms or molecules packing in regular or repeating way all the way through the solid
44
define cation
any ion with an overall positive charge
45
define amorphous solid
not ordered and molecules are packed at random
46
define anion
any ion with an overall negative charge
47
define lattice enthalpy
the energy change which occurs when one mole of solid forms from ions which start out indefinitely far apart, and since the lattice forms and makes bonds, the system ends up lower in energy and so we represent that as a negative enthalpy
48
define electronegativity
how much an element draws electrons towards itself
49
define the Octet rule
each atom tries to adopt a state with 8 electrons in its highest energy level
50
define the Bond-Oppenheimer approximation
electrons are small and can move quickly, whilst atomic nuclei are larger, heavier and do not move as fast, so we can regard the nuclei as static on the timeframe of electrons
51
define quantum confinement
the phenomenon by which the more localised a particle is, the higher energy it is
52
what are the limitations of the particle in a box model
atoms are not one dimensional, the electrostatic potential well for a Coulombic potential is not boxed shaped, and the atomic potential well is not infinitely deep as you are able to remove electrons from atoms
53
define valence bond theory
bonds form when two atoms orbitals overlap and share two electrons between themselves
54
how do we get in-phase combination
by adding two orbitals together
55
how do we get out-of-phase combination
by subtracting one orbital from the other
56
what does the in-phase combination show
constructive interaction
57
What is the difference between in-phase and out-of-phase combination
in-phase: the space occupied by the electrons in the combined orbital is larger than the space occupied by the electrons in the two separate orbitals
out-of-phase: the space occupied by the electron is less than that of the free atoms
58
define bonding orbital
the in-phase combination which produces an orbital lower in energy than the original orbitals
59
what is the equation for bond order
1/2 (number of electrons in bonding orbitals - number of electrons in anti-bonding orbitals)
59
isoelectric relationship
two compounds with different atoms and charges have the same number of electrons in the same type of orbital
59
define anti-bonding orbital
an out-of-phase combination which is higher in energy than the original orbitals
60
what are the rules for filling MO diagrams
fill from the lowest energy up
for orbitals of the same energy, electrons fill up first with spins parallel, then anti-parallel pairing
61
define orthogonality
two orbitals can occupy the same space but never form bonding or anti-bonding combinations with each other
62
define sigma symmetry
if you rotate around the axis of the atom then everything stays the same
63
define pi symmetry
if rotate around the axis and the wave function changes
64
define inversion symmetry
if (+x, +y, +z) and (-x, -y, -z) are equal and opposite
65
what is the label we give something if it has inversion symmetry
u, meaning ungerade
66
what label do we give something if πΉ(+π₯, +π¦, +π§) = πΉ(βπ₯, βπ¦, βπ§)
g, for gerade
67
define paramagnetic
compounds with unpaired electrons
68
define diamagentic
compounds with paired electrons
69
define s-p mixing
when the s and p orbitals on adjacent atoms can interact if they point towards each other
70
define molecular orbital theory
the true MOs of a molecule are formed by LACO (linear combination of AOs) and are determined by considering all of the AOs at once rather than only considering pair-wise interactions
71
define local orbital approximation
the idea that only nearby things will significantly affect the local environment
72
define conjunction
used to describe the situation when pi systems are 'linked together'
73
define hybridization
the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory
74
describe how to determine hybridization of an atom in a molecule
1. Count the number of atoms (not bonds) connected to it
2. Count the number of lone pairs attached to it
3. Add these numbers together
4. If itβs 4, the atom is sp3
5. If itβs 3, the atom is sp2
6. If itβs 2, the atom is sp
7. If itβs 1, itβs probably hydrogen
75
define resonance stabilisation
the kind of stabilisation through delocalisation into conjugated systems
76
why is the kekule structure incorrect
- It predicts alternating short and long bonds, but all measurements are actually the same
- Double bonds react rapidly with Br2 as this is one of the common tests for alkenes, however benzene does not
- It predicts that 1,2-disubstituted benzene should show two isomers, but benzene only shows one
- Benzene is much more stable, thermodynamically speaking, than would be expected. It would be expected to be -406kJmol-1, but in reality it is -206kJmol-1
77
what relationship must be true for a compound to be aromatic
ππ. ππ π πππππ‘ππππ = 4π + 2 (Huckel's rule)
78
what is the relationship for anti-aromatic compounds
they have 4n electrons
79
how does photoelectron spectroscopy work
by ionizing samples using high energy radiation (UV or X-rays) and measuring the kinetic energies of the ejected electrons
80
what does photoelectron spectrsocopy determine
the relative energies of electrons in atoms or molecules
81
