AP Chemistry
Winter Break Assignment
Name
Date
This assignment will encompass the concepts concerning electronic structure, Bohr model
of the atom, quantum numbers, orbital shapes, sizes, and hybridization, the Quantum
Mechanical model of the atom, electron configuration, periodic trends, lewis structures,
molecular geometry and polarity, phase diagrams, intermolecular forces, and crystal
structure of molecular and ionic solids. You should know and understand the concepts
behind these topics upon the return to school on January 2, 2013. Calculations will be
reviewed before a graded assessment is assigned, but the basic concepts will not be
addressed in depth and it is your responsibility to study the material provided, along with
reading Chapters 6, 7, and 9 of the text. Problem sets listed below will be due on Tuesday,
January 8, 2013. Complete all worksheets provided and attach the problem set to this
packet.
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Chapter 6: page 158- Summary Problem, #1, 2, 3, 6, 7, 8, 10, 11, 14, 19, 27, 29, 31,
39, 50, 61, 67, 71, 73, 76, 80
Chapter 7: page 191-Summary Problem, #17, 19, 21, 24, 27, 29, 37, 47, 55, 57, 63,
69, 81
Chapter 9: page 252- Summary Problem, #1, 3, 5, 7, 9, 11, 17, 21, 31, 33, 45, 51, 56,
60, 61, 66, 74, 76
Chapter 6: Electronic Structure and the Periodic Table
6.1: Light, Photon Energies, and Atomic Spectra
 Structure of the Atom: the structure of the atom has changed drastically over time,
beginning with Democritus’ idea of the atom being an indivisible sphere and
continuing to evolve today.
o 450 B.C.
 Democritus
 All matter is made up of “atomos” (indivisible particles)
 At a certain point, matter could not be broken down any
further
 Aristotle
 Disagreed with Democritus
 4 Elements: Earth, Wind, Fire, Water
 Both were philosophers, not scientists. Aristotle’s ideas were believed
for nearly 2000 years.
o Late 1700s: 3 important discoveries
 Antoine Lavoisier: Law of Conservation of Matter
 Joseph Proust: Law of Definite Proportions
 Ben Franklin: Electricity
 Objects could have 1 of 2 electrical charges (+/-)
 Like charges repel, opposites attract
 Some objects readily pick up electrical charge (static
electricity, metals)
o 1803
 John Dalton and his Atomic Theory
 Each element is composed of small particles (atoms)
 All atoms of the same element are identical, but differ from
those of any other element
 Atoms are neither created or destroyed in a chemical reaction
 A given compound always has the same relative number and
kinds of atoms
 Atoms are small, solid, spherical, indivisible, neutral
 MARBLE MODEL
o 1840
 Michael Faraday
 Atomic Structure is related to Electricity!
 Used Cathode Ray Tubes and discovered a stream of particles
 **Cathode Ray tube is a partially evacuated glass tube
containing a gas at a low pressure**
o 1869
 Dmitri Mendeleev and Lothar Meyer published almost identical
periodic tables. (Mendeleev got more credit than Meyer)
o 1871
 Mendeleev predicted unknown elements and their properties.

o 1896
 JJ Thomson
 Used Cathode Ray Tube to study electric discharge in a vacuum
 Used A NODE with a hole to allow a ray to pass through
 Magnets and Electricity can alter the ray’s path in a
mathematically predictable way
 Concluded the cathode ray was NEGATIVE PARTICLES coming
from the cathode (ELECTRONS)
 **Proved Democritus wrong (Atoms are divisible)
 Charge to mass ratio = 1.76 x 108 C/g
 PLUM PUDDING MODEL/CHOCOLATE CHIP COOKIE MODEL
 Henri Becquerel
 Radioactivity: Spontaneous emission of radiation
 Marie Curie and Pierre Curie followed in his footsteps
o 1900
 MAX PLANCK
 Proposed that there was a predictable relationship between a
QUANTUM of energy and the frequency of radiation
 An object emits energy in small, specific amounts (QUANTA)
 E = h
h = Planck’s Constant = 6.626 x 10-34Js
 Planck’s suggestions were thought to be extremely radical at
the time
1905
o ALBERT EINSTEIN
 Expanded on Planck’s theory by stating that electromagnetic radiation
(EMR) has a dual wave-particle nature
 Each individual particle of light can carry 1 quantum of energy
 PHOTON: a particle of EMR having zero mass and carrying a quantum
of energy
 Ephoton = h
 Minimum energy ~Minimum frequency
 Must reach this minimum energy before a photon will be released
o 1909
 Robert Millikan
 Oil Drop Experiment: Spray drops in apparatus, give drops
negative charge using x-rays
 Varying the charge effected droplets’ rate of fall
 Determined charge of electron: 1.60 x 10-19C
 Mass of electron = 9.11 x 10-28g
o 1911
 Ernest Rutherford
 Gold Foil Experiment: Bombarded a thin sheet of gold with
alpha particles and most of them passed through. A few
bounced back
 Discovery of NUCLEUS


Suggested that electrons orbit nucleus, but could not explain
his findings.
 Nucleus: Densely packed bundle of matter with a positive
charge
Henry Moseley
 Elements fit into better patterns when arranged by increasing
ATOMIC NUMBER (# of protons)
 PERIODIC LAW: The physical and chemical properties of the
elements are periodic functions of their atomic numbers.
 PERIODIC TABLE: An arrangement of the elements in order of
atomic number so that elements with similar properties
(PERIODICITY) are in the same GROUP.
 PERIODICITY: properties repeat themselves at regular
intervals throughout the periodic table.
o 1913
 Niels Bohr
 He developed a model of the atom that linked ELECTRONS
with PHOTON EMISSION
 Electrons can circle the nucleus in specific paths (ORBITS)
and when in one of these orbits, the atom has a DEFINITE
and FIXED ENERGY. Within each orbit, their exists
SUBLEVELS.
o Orbit closest to the nucleus = lower energy state
o Orbits further from nucleus = higher energy state
o Electrons CANNOT EXIST between orbits (energy
levels)
 PLANETARY/SOLAR SYSTEM MODEL
*Bohr’s idea explained only Hydrogen, not atoms with more than 1 electron*
o QUANTUM MECHANICAL MODEL: Differs from the Bohr Model in that
 The Kinetic Energy of an electron is inversely related to the volume of
the region to which it is confined
It is impossible to specify the precise position of an electron in an
atom at a given instant
o 1924: LOUIS DE BROGLIE
 Bohr’s quantized electrons behaved like waves
 Duality of matter: Electrons could be considered as waves confined to
a certain space around an atomic nucleus. (certain space = specific
frequency = specific energy)
 Experimentation proved that:
1. Electrons (like light waves) could be bent of DIFFRACTED
a. Diffraction: The bending of a wave as it passes by the edge of an
object (ex: edge of an atom in a crystal)

2. Electron beams (like waves) can INTERFERE with each other
a. Interference: Waves overlap which results in a reduction of energy
in some areas and an increase of energy in other areas
o HEISENBERG UNCERTAINTY PRINCIPLE
 Detection of electrons by observing their interaction with photons
 Photons have about the same energy as electrons and when used to
detect electrons the photon will change the electron’s path. This
results in UNCERTAINTY when trying to locate electrons
HEISENBERG UNCERTAINTY PRINCIPLE states: it is impossible to
determine simultaneously both the POSITION and VELOCITY of an
electron (or any particle)
 We must rely on PROBABILITY that an electron is in a certain location
and moving at a certain speed.
 Photons are shot at an electron and when they collide we are able to
find the exact LOCATION of an electron, but this collision causes the
VELOCITY to change.
o SCHRODINGER WAVE EQUATION
 Schrodinger used the dual wave-particle hypothesis to develop an
equation that treated electrons in atoms as waves
 Naturally accounted for Bohr’s idea of quantization
 Together with HUP, laid a foundation for modern QUANTUM THEORY:
Describes mathematically the wave properties of electrons and other
small particles
 Based on HUP, we can only find the probability of locating an electron
at a given point around the nucleus
 Electrons did not exist in neat paths (BOHR’S ORBITS) but instead in
certain regions (ORBITALS: 3-D region around the nucleus that
indicates the probable location of an electron)

o PAULI EXCLUSION PRINCIPLE: No 2 electrons can be in the same exact place
at the same exact time.

ELECTROMAGNETIC SPECTRUM
 Made up of all forms of ELECTROMAGNETIC RADIATION
 All move at a constant speed of 2.998 x 108 m/s (167,000,000 mi/hr)
WAVELENGTH: ( : lambda) The distance between corresponding points on
adjacent waves
o 1 nm = 1 x 10-9m
o This value must always be converted to METERS
 FREQUENCY: ( : nu) The number of waves that pass a given point in a specific
time
o 1 wave/second = 1Hertz (Hz)
 c = 
 Amplitude: ( Ψ: psi) Height of crest/trough
Photon Energies
o Light is a stream of tiny particles (photons) and the energy can be calculated
using the following equation
o E = h
 Energy is measured in Joules
 h = Planck’s Constant = 6.626 x 10-34Js
Atomic Spectra: All species emit photons at specific, predictable wavelengths. These
wavelengths (and their corresponding frequencies and energies) can be studied to
determine the identity of the species.
o The simplest of the atomic spectra is HYDROGEN. Hydrogen’s atomic spectra
has been studied extensively in an effort to understand more about
electronic behavior. Hydrogen emits groups of photons in 3 different regions
of the electromagnetic spectrum:
 Lyman Series (UV region)
 Balmer Series (Visible region)
 Paschen Series (Infrared region)
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
6.3: Quantum Numbers
1. PRINCIPAL QUANTUM NUMBER (n)
a. Main energy level of electron
b. Positive whole numbers (1,2,3,…)
c. Energy increases as distance from the nucleus increases
d. More than 1 electron can have the same n-value
Shapes of orbitals:
e. Total number of orbitals in a given shell = n2
2. ANGULAR MOMENTUM QUANTUM NUMBER (l)
a. Indicates the sublevel, which determines the general shape
of the electron cloud
b. This value ranges from 0 – (n-1)
c. l = 0 = s-sublevel
1 = p-sublevel
2 = d-sublevel
3 = f-sublevel
ENERGY
LEVEL
1
SUBLEVEL
# of ORBITALS
# of ELECTRONS
s
1
2
2
p
3
6
3
d
5
10
4
f
7
14
3. MAGNETIC QUANTUM NUMBER (ml)
a. Designates the direction in space in which the orbital is
oriented
b. More than 1 electron can have the same SHAPE, but a
different ORIENTATION
c. –l to +l
4. SPIN QUANTUM NUMBER (ms)
a. TWO possible values (+/- ½)
b. TWO fundamental spin states of an electron in an orbital
c. A single orbital can hold a maximum of 2 electrons with
opposing spins
6.5: Electron Configurations in Atoms

ELECTRON CONFIGURATIONS
o The arrangement of electrons in an atom
o Distinct address for each element
o Electrons will assume arrangements with lowest
o possible energy
o RULES:
1. ORDER of electrons occupying each orbital
a. Aufbau Principle: an electron occupies the lowest energy
level/orbital that can receive it
2. Importance of SPIN
a. Pauli Exclusion Principle: no 2 electrons in the same atom
can have the same set of 4 QNs
3. Electrons will occupy ALL orbitals before pairing in an orbital
a. Hund’s Rule: orbitals of equal energy are each occupied by
a second electron and all electrons in singly occupied
orbitals must have the same spin
o 3 METHODS OF NOTATION
1. Orbital Notation
a. Unoccupied orbital represented by a line:
b. Orbital containing one electron: 
c. Orbital containing 2 electrons:  (opposing spins)
d. Each line (orbital) is labeled with Principal Quantum
Number
2. Electron-Configuration Notation
a. Eliminates lines and arrows of orbital notation
b. Electrons in each sublevel indicated by adding a
SUPERSCRIPT to the sublevel
3. Noble-Gas Notation
a. Beginning with 3rd period, electron-configuration notation
can be abbreviated using the noble gases
6.8: Periodic Trends
1. ATOMIC RADIUS
 ½ the distance between the nuclei of 2 identical atoms that are
bonded together
 DECREASES from left to right
 INCREASES from top to bottom


2. IONIC RADIUS
 Formation of a CATION always leads to a DECREASE in atomic radius
 Formation of an ANION always leads to an INCREASE in atomic radius
3. IONIZATION ENERGY
 An ion is an element/group of elements that has a positive or negative
charge
 Ionization is any process that results in the formation of an ion
 The energy required to remove one electron from a neutral atom of an
element.
 Those elements that EASILY lose electrons are HIGHLY REACTIVE
 INCREASES from left to right (due to increasing nuclear charge)
 DECREASES from top to bottom (due to increasing size of electron
cloud)
 Electrons further from the nucleus are more easily removed
 SUCCESSIVE IONIZATION ENERGY: Removal of electrons from
cations. Each successive electron removed from an ion feels an
increasingly stronger effective nuclear charge.
 Noble Gases and Ions with Noble Gas Electron Configuration have
extreme stability making it very difficult to lose more electrons

4. ELECTRONEGATIVITY
 A measure of the ability of an atom in a chemical compound to
ATTRACT electrons
 INCREASES from left to right
 DECREASES from top to bottom

5. ELECTRON AFFINITY
 Energy change that occurs when a neutral atom acquires an electron
 INCREASING or DECREASING (no real pattern)
 DECREASES from top to bottom

Chapter 7: Covalent Bonding
7.1: Lewis Structures and the Octet Rule



1916: Gilbert Newton Lewis suggested that noble gases are very unreactive and that
nonmetal ions can achieve this stability by sharing electrons.
The Octet Rule states that the most stable arrangement of electrons consists of 8
valence electrons, resembling the noble gases.
o Exceptions to the Octet Rule
 Hydrogen and Helium
 Boron: generally forms 3 bonds because it has only 3 valence
electrons
 Elements bonded with F/O/Cl (Highly Electronegative) sometimes
hold more than 8 valence electrons; an expanded octet or expanded
valence shell
Lewis Structures
o Atomic Symbols represent nuclei and inner-shell electrons.
o Dots/Dashes between atoms represent SHARED PAIR of electrons.
o Dots next to only one atom represent LONE PAIR of electrons.
o Structural Formula = only shared pairs between atoms are represented,
usually by dashes but can also be represented by lines.
o RULES
1. Count the # of valence electrons
2. Write symbols of LEAST ELECTRONEGATIVE elements in the MIDDLE
3. Assign dots as valence electrons around all elements. Make sure not to
exceed the total number of valence electrons.
4. If each element in the structure does NOT have 8 valence electrons,
rearrange to form double or triple bonds. If there are electrons left over,
add electron pairs to central atom. (This usually only happens when the
central atom is bonded to HIGHLY ELECTRONEGATIVE elements)
o Multiple Covalent Bonds
 Double bond: 2 shared pairs between 2 atoms.
 Triple bond: 3 shared pairs between 2 atoms.
o Resonance Structures: Bonding cannot be accurately represented by a single
Lewis Structure
Ex: O3 (ozone)
O = 6(3) = 18ve
o Ions: atom/group of atoms that has a charge
 (+) ion (cation): subtract charge from total number of valence
electrons
 (-) ion (anion): add charge from total number of valence electrons
7.2: Molecular Geometry

VSEPR Theory: Valence-shell electron-pair repulsion theory
o Electrons repel each other, causing bonds to be as far away from one another
as possible.
o Most important atom when determining geometry is the CENTRAL ATOM
o ABE Notation: describes how electrons are shared around the central atom.
 A = Central Atom
 B = # of atoms bonded to central atom
 E = # of lone pairs on central atom
SHAPE
Linear
Bent/Angular
Trigonal Planar
Tetrahedral
Trigonal Pyramidal
Trigonal
Bipyramidal
Octahedral
# ATOMS BONDED
TO CENTRAL (B)
1
2
2
2
3
4
3
5
# LONE PAIRS ON
CENTRAL (E)
3,1,0
0
1
2
0
0
1
0
ABE NOTATION
ABE3/ABE/ABE2/AB
AB2
AB2E
AB2E2
AB3
AB4
AB3E
AB5
6
0
AB6
7.3: Polarity of Molecules
 Intermolecular Forces (IMF): Forces of attraction between MOLECULES
o Weaker than bonds (covalent/ionic/metallic) holding atoms together.
o As temperature increases, KE increases, and molecules begin to pull away
from one another.
o IMFs are between MOLECULES (Covalently bonded groups of atoms)
o 3 Types:
1. London Dispersion Forces
 Results from the constant motion of electrons and creates
instantaneous dipoles
 Present in ALL ATOMS and MOLECULES
 Noble Gases and Nonpolar Molecules only have this type of IMF
 Increase with increasing # of electrons and increasing mass
2. Dipole-Dipole Forces
 Dipole: created by equal but opposite charges separated by a
distance (bond length)
 Ex: H2O
  = Delta
 + = dipole; (+) near less electronegative element,  near more
electronegative element
 Forces of attraction between POLAR MOLECULES
 POLARITY depends on both ELECTRONEGATIVITY difference as
well as GEOMETRY
GEOMETRY
ALWAYS NONPOLAR
ALWAYS POLAR
SOMETIMES POLAR

Linear

Bent/Angular

Trigonal Planar

Tetrahedral

Trigonal Pyramidal

Trigonal
Bipyramidal

Octrahedral
 Dipole-Induced Dipole: A polar molecule can INDUCE a dipole in a
nearby NONPOLAR molecule by temporarily attracting its electrons
(This force is WEAKER than a Dipole-Dipole force)
3. Hydrogen Bonding
 Hydrogen is bonded to the following highly electronegative
elements
i. NITROGEN
ii. OXYGEN
iii. FLOURINE
 Strongest IMF
 Ex: H2O
7.4: Atomic Orbitals and Hybridization

sp hybrid orbitals: Bonding electrons located in the s-orbital of 1 atom and the porbital of a second atom will for 2 sp-hybrid orbitals
o s + p  2sp

sp2 hybrid orbitals: s + 2p  3sp2

sp3 hybrid orbitals: 1s + 3p  4sp3

sp3d hybrid orbitals: 1s + 3p + 1d  5sp3d

sp3d2 hybrid orbitals: 1s + 3p + 2d  6sp3d2

Extra electrons in a multiple bond are NOT located in hybrid orbitals. Only 1 of the
shared pairs in a multiple bond exists in the hybrid orbital.
o Sigma Bonds ( σ ): Consist of an electron pair occupying a hybrid orbital
o Pi Bonds ( π ): Consist of an electron pair occupying the unhybridized orbital
 SINGLE BOND: 1 sigma
 DOUBLE BOND: 1 sigma and 1 pi
 TRIPLE BOND: 1 sigma and 2 pi
Chapter 9: Liquids and Solids
9.2: Phase Diagrams
9.4: Network Covalent, Ionic, and Metallic Solids
1. Network Covalent Solids: a continuous network of covalent bonds joins atoms.



High Melting Point
Insoluble in all common solvents
Poor electrical conductors
2. Ionic Solids: Consist of cations and anions




Nonvolatile with high melting points
Do not conduct electricity unless dissolved or in the liquid state because as solids,
ions are in fixed, permanent positions
Many are soluble in water
Strength of ionic bonds can be estimated using Coulomb’s Law and depends on 2
factors
o Charges of the ions
o Size of the ions
3. Metals: Consist of a sea of electrons






High electrical conductivity
High thermal conductivity
Ductility and Malleability
Luster
Insolubility in water and other common solvents
Mercury can dissolve other metals to form a solution called an amalgam.
9.5: Crystal Structures
**Table 9-5: Structures and Properties of Types of Substances (p.245)**
A unit cell is used to represent the structural geometry of a solid
 Metals: 3 types of solid structures
o Simple Cubic Cell (SC): 8 atoms
o Face-Centered Cubic Cell (FCC): 9 atoms
o Body-Centered Cubic Cell (BCC): 12 atoms
**Table 9-6: Properties of Cubic Unit Cells**

Ionic Crystals: 3 types of lattices (see page 249 for diagrams of ionic crystals)
o Face-Centered Cubic Lattice-1 (FCC1)
o Face-Centered Cubic Lattice-2 (FCC2)
o Body-Centered Cubic Lattice (BCC)
Wavelength, Frequency, and Energy Calculations
1. Green light has a wavelength of 5.0 x 102 nm.
a. What is the energy, in joules, of one photon of green light?
b. What is the energy, in joules, of 1.0 mol of photons of green light?
2. Violet light has a wavelength of about 410 nm.
a. What is the frequency?
b. What is the energy of one photon of violet light?
c. What is the energy of 1.0 mol of violet light?
3. The most prominent line in the spectrum of aluminum is 396.15 nm.
a. What is the frequency of this line?
b. What is the energy of one photon with this wavelength?
4. The most prominent line in the spectrum of magnesium is 285.2 nm. Other lines are
found at 383.8 nm and 518.4 nm.
a. In what region of the electromagnetic spectrum are these lines found?
b. Which is the most energetic line?
c. What is the energy of 1.00 mol of photons with the wavelength of the MOST
ENERGETIC line?
5. The light from an amber signal has a wavelength of 595 nm, and that from a green
signal has a wavelength of 500.0 nm.
a. Which has a higher frequency?
b. Calculate the frequency of amber light.
6. You are an engineer designing a switch that works by the photoelectric effect. The
metal you wish to use in your device requires 6.70 x 10-19J/atom to remove the
energy from a single electron. Will the switch work if the light falling on the metal
has a wavelength of 540nm or greater? Why or why not?
7. If a large pickle is attached to two electrodes with a 110-Volt power supply, the
pickle begins to glow with a yellow color. Knowing that pickles are made by soaking
cucumbers in a concentrated salt (NaCl) solution describe why the pickle might emit
yellow light when electrical energy is added.
Quantum Numbers and Electron Configuration
1. Determine the 4 quantum numbers for the following elements
a. N
b. Mg
c. Ar
d. Fe
e. Rb
f. Br
2. What are 2 exceptions to electron configurations and the Aufbau Principle, and why
do these exceptions occur?
3. What are the 4 Quantum Numbers for the following electrons?
a. 4d3
b. 6s1
c. 2p5
4. For each of the following elements, write the Orbital Notation, Electron
Configuration Notation, and Noble Gas Notation on a separate sheet of paper.
a. Li
b. Mg
c. G
d. C
e. P
f. Se
g. F
h. Kr
i. Cr
j. Ag
5.
6.
7.
8.
9.
When n = 4, what are the possible values of l?
When l = 2, what are the possible values for ml?
In a 4s orbital, what are the possible values of n, l, ml and ms?
When n = 4, l = 2, and ml = -1, what orbital does this refer to? (Ex: 1s)
Why is the following set of quantum numbers not possible for an electron in an
atom?
n = 2, l = 2, ml = 0
10. Which of the following orbitals cannot exist according to the quantum theory: 2s, 2d,
3p, 3f, 4f, and 5s. Explain your answer.
11. A given orbital has a magnetic quantum number of ml = -1. Which sublevel is not
possible? (Ex: s, p, d, or f)
Periodic Trends
1. Which has a larger atomic radius: Sodium or Cesium? Explain your reasoning.
2. Which has a smaller radius: Calcium atom or Calcium +2 ion? Explain your
reasoning.
3. Which has a higher 1st ionization energy: Lithium or Oxygen? Explain your
reasoning.
4. Which ionization energy for Radium is the highest and explain your answer:
a. 1st
b. 2nd
c. 3rd
d. 4th
5. Which has a lower electronegativity value: Silicon or Sulfur? Explain your
reasoning.
1. Which has a larger atomic radius: Aluminum or Silicon
2. Which has a smaller atomic radius: Strontium or Rubidium
3. Which has a larger ionic radius: Fluorine atom or Fluorine ion
4. Which has a smaller ionic radius: Calcium atom or Calcium ion
5. Which has a larger 1st ionization energy: Lithium or Neon
6. Which has a larger 2nd ionization energy: Phosphorus or Sulfur
7. Which has a higher electronegativity: Chlorine or Iodine
8. Which has a lower electrongativity: Magnesium or Barium
9. Which ionization energy for Gallium is the highest: 1st, 2nd, 3rd, or 4th
Lewis Structures, Molecular Geometry, Polarity, and Intermolecular Forces
DIRECTIONS: For each of the following compounds,
a) Draw Lewis Structures (be sure to indicate the number of valence electrons)
b) Identify the ABE notation
c) Predict the geometry for each of the molecules
d) Predict the polarity of the molecule (If polar, show direction of polarity)
e) State the strongest IMF present
1. AsCl5
2. SeF6
3. PH3
4. SF2
5. C2HI
6. BCl3
7. CH3CHCCHCH3
8. HOCl
9. CHF3
10. PCl5
11. SeF6
12. PH3
13. Cl2CO
14. SiO2
15. CH3CH2OCH3
16. CH3CH2CH2CH2NH2
17. HI
18. H2Se
19. OClOH
20. CH4
21. (CO3)2—
22. NH4+
23. C2H6
24. CO2
25. HCN
26. C3H8
27. C3H6
28. BH3
29. SeH6
30. F2
31. CH3COOH