Name:
Date:
Chemistry II Notes – Atoms, Molecules, and Ions
HISTORY OF ATOMIC THEORY:
 The Greek philosopher Democritus (460-370 BC) was the first person to propose the idea of matter being
composed of small, indivisible particles (atomos)
 In the late 18th century, Lavoisier proposes the law of conservation of mass & Proust proposes the law of
definite proportions
 In 1808, Dalton proposed the atomic theory using these previously unconnected ideas
ATOMIC THEORY:
 DALTON’S ATOMIC THEORY (MODERN):
o All matter is composed of atoms.
o An atom is the smallest particle of an element that maintains the properties of that element. (atoms
defined)
o Atoms of one element are different in physical and chemical properties from atoms of other
elements.
o Atoms of different elements combine in simple whole-number ratios to form compounds. A given
compound always has the same relative numbers and types of atoms (Law of definite proportions).
o Atoms are separated, combined, or rearranged in chemical reactions – they cannot be created nor
destroyed (Law of conservation of mass).
ATOMIC STRUCTURE:
 Several experiments were being carried out in the 19th and 20th centuries that began to identify the subatomic particles that make up the atom. A summary of those experiments is given below:
Scientist:
Experiment:
Crookes
Cathode Ray Tube
J. J. Thomson
Cathode Ray Deflection
Millikan
Oil Drop Experiment
Rutherford, Marsden, &
Geiger
Gold Foil Experiment
Knowledge
Gained:
Negative particles of some kind
exist
Mass/charge ratio of electron
determined (1.76 x 108 C/g)
Charge on the electron
(1.60 x 10-19 C)
Nucleus present in the atom
Relating To…
Electron
Electron
Electron
The nucleus of an
atom and the proton
*Knowing the charge-to-mass ratio of the electron & the charge of the electron, the mass is then calculated to be 9.10 x 10-28g*


In the first part of the 20th century, following Chadwick’s discovery of the neutron, Bohr proposed the idea
that the atom was made up of the nucleus containing protons and neutrons that was being orbited by
electrons in specific, allowed orbits.
This particle model of the electron and atom was expanded a few years after Bohr’s original ideas to
incorporate the wave nature of the electrons. (More about this later…)
Particle:
Proton
Charge:
+1
Mass:
1.0073 amu
Position in Atom:
Nucleus
Neutron
0
1.0087 amu
Nucleus
Electron
-1
5.486 x 10-4 amu
Outside of the nucleus
Misc.:
Identify the element; composed of 2 ‘up’
& 1 ‘down’ quarks
Form isotopes for an element; composed
of 2 ‘down’ & 1 ‘up’ quarks
Affect reactivity of an element
*amu = atomic mass units; 1 amu = 1.66054 x 10-24g*



Most of the mass of the atom is due to the nucleus.
Most of the volume of the atom is the space, outside of the nucleus, where electrons are found.
The quantity 1.602 x 10-19 C is termed the electronic charge and it is equal to the charge on an electron and
the charge on a proton.
1
Name:






Date:
Since masses are so small, we define the atomic mass unit (amu)
The angstrom is a convenient non-SI unit of length used to denote atomic dimensions.
o Since most atoms have radii around 1 x 10-10 m, we define 1Å = 1 x 10-10 m
Atomic number (Z) = number of protons in the nucleus
o Since all atoms are neutral it also tells us the number of electrons surrounding the nucleus
 When atoms lose or gain electrons they become charged and form ions
Mass number (A) = total number of nucleons in the nucleus (protons & neutrons)
o On most periodic tables, atomic mass numbers are not integers. This indicates that there is more
than one isotope of the element existing in nature.
 Example: On many periodic tables the atomic mass of Cl is listed as 35.5. This does NOT
indicate that there are 17 protons, 17 electrons, and 18.5 neutrons in an atom of
chlorine. It is not possible to have a fraction of a neutron, there can only be a whole
number of neutrons in an atom. The non integer values mean that there is more than one
isotope of chlorine that exists in nature, in this case 35Cl and 37Cl. A quick calculation
will tell you that these two species have the same number of protons and electrons, but
different (whole) numbers of neutrons (18 and 20 respectively). So, they are isotopes of
one another. These isotopes happen to exist naturally in the following abundance: 35Cl
75% and 37Cl 25%.
By convention, for element X, we write A to Z notation
o Atomic mass (A) is written as a superscript
o Atomic number (Z) is written as a subscript
Isotopes have the same Z (protons/electrons) but different A (number of neutrons)
o Since it is the electrons that affect the chemical properties of a substance, isotopes of the same
element have the same chemical properties.
o There can be a variable number of neutrons for the same number of protons. Isotopes have the
same number of protons but different numbers of neutrons
o An atom of a specific isotope is termed a nuclide
 Example: Nuclides of hydrogen include: 1H = hydrogen (protium); 2H = deuterium;
3
H = tritium (tritium is radioactive)
PERIODIC TABLE:
 The periodic table is used to organize the elements is a meaningful way. So, there are periodic properties
associated with periodic table arrangement.
o Columns = groups/families
o Rows = periods
 Periodic law: When elements are arranged by increasing atomic number, there is a periodic repetition of
their physical and chemical properties
o Elements in the same groups have the same number of valence electrons
 This similarity gives them similar properties
 Differences arise from their electrons being in different shells
 Major sections of the periodic table: metals & nonmetals
o METALS tend to be malleable, ductile, and lustrous and are good thermal and electrical
conductors. Metals are located on the left-hand side of the periodic table (most elements are
metals) and are typically solid at room temperature (exception of note is mercury – liquid at room
temperature). Metals tend to lose electrons to form positive ions.
o NONMETALS generally lack these properties; they tend to be brittle as solids, dull in appearance,
and do not conduct heat or electricity well. Nonmetals are located in the top right-hand side of the
periodic table (exception of note is hydrogen – upper left-hand corner). Nonmetals tend to gain
electrons to form negative ions; often bond to each other by forming covalent bonds; and typically
are gases at room temperature.
 Minor sections: metalloids & noble gases
o METALLOIDS have properties similar to both metals and nonmetals and are located at the
interface between the metals and nonmetals (stair steps). Metalloids tend to be less malleable than
metals/not as brittle as nonmetals, semiconductors, and solids at room temperature. Metalloids are
considered to be the following elements: B, Si, Ge, As, Sb, Te, Po, At.
2
Name:
Date:
o
NOBLE GASES are generally unreactive and under normal conditions exist as monatomic
(single-atom) gases. Noble gases are located on the far right-hand side of the periodic table (group
18)
RADIOACTIVITY:
 Radioactivity is the spontaneous decay of certain atoms with the evolution of alpha, beta, and gamma
particles.
 The radiation comes from the nucleus of the atom (it is a nuclear reaction)
Nature:
Alpha α
A helium nucleus
Beta β
Essentially electrons
Charge:
+2
-1
Gamma γ
High energy & high
frequency EMR
0
Mass:
4
5.486 x 10-4
0
Movement in Electric
Field:
Penetrating Power:
Toward negative plate
Toward positive plate
None
Least
Intermediate
Greatest





RADIOACTIVE DECAY REACTIONS
o Alpha decay (the loss of a Helium nucleus)
 mXn → m-4Yn-2 + 4α2
o Beta decay (a neutron splits to give a proton and an electron)
 mXn → mZn+1 + 0β-1
o Gamma decay (a rearrangement of the nuclear particles only)
 mXn → mXn + γ (gamma radiation)
OTHER NUCLEAR REACTIONS
o Positron emission (a positron has the same mass as an electron, but it has a positive charge)
 mXn → mZn-1 + 0β+1
o Electron capture (a captured electron combines with a proton in the nucleus to form a neutron)
 mXn + 0e-1 → mZn-1
Half-life of a radioactive nucleus is the time taken for half of the atoms to decay. It is independent of the
initial quantity of atoms and there are three methods of determining half-life:
o Graphically
o Use of the expression: 2.303log(N0/N) = kt
 Where N0 = count rate initially; N = count rate at time t; k = radioactive constant; and t =
time.
 If the half-life is known then N0/N = 2 and 2.303log2 = kt1/2
o Use of the expression: fraction of remaining activity = 1/2 n where n = number of half-lives
Transmutation of elements - it is possible via nuclear reactions, to artificially produce elements.
o 1919 Rutherford – Alpha particle bombardment
 14N7 + 4He2 → 1H1 + 17O8
o 1932 Cockcroft and Walton. H+ used as the bombarding particle, it has less dense charge than the
alpha particles and so feels less repulsion from the target nucleus.
 7Li3 + 1H1 → 4He2 + 4He2
o Neutron induced transmutation. Neutrons being neutral feel no repulsion from the positive target
nucleus.
 31P15 + 1n0 → 32P15 + gamma
o 1970 Use of accelerated heavier nuclei. A machine will accelerate the positive nucleus toward the
target.
 250Cf98 + 11B5 → 257Lw103 + 41n0
Mass deficit – when atoms are formed by the combination of protons, neutrons and electrons the mass of
the atom is found to be less than that of the sum of the individual particles.
o This appears to contradict the law of conservation of mass
3
Name:
Date:
o





The explanation for the mass deficit is that when the particles combine a small amount of the mass
is converted to energy (binding energy) and released to the surroundings
Predicting stability – stable (non-radioactive) nuclei tend to have neutron:proton ratios close to 1:1. Nuclei
that have higher neutron:proton ratios tend to want to lower the ratio by converting a neutron to a proton
and an electron. The electron is released as β particles.
Nuclear and electron arrangement – radioactivity is a nuclear process that involves rearrangement of the
nuclei of atoms, whereas chemical reactions involve the rearrangement of the electrons
Nuclear fission – the process of heavy nuclei capturing neutrons, splitting to form other, smaller nuclei and
releasing more neutrons. In the process, large amounts of energy can be released, exacerbated by the
production of more neutrons each time leading to a potential chain reaction.
Nuclear fusion – the combination of smaller nuclei into larger ones with the release of energy. These
reactions are less easy to perform than nuclear fission since they involve the combination of two nuclei that
are both positively charged and therefore repel one another.
Uses of radioactivity – some examples are listed below:
o Medicine – such as 131I for thyroid and brain imaging, 58Co for diagnosis of pernicious anemia,
and 67Ga for lung function
o Isotopic dating
o Thickness control in engineering
 A radioactive source is placed on one side of a sheet of paper or metal as it is made and a
detector is placed on the other side.
 Any change in the thickness of the material will be reflected in a change in the observed
radioactivity count.
o Leak detection
 Radioactive sources can be injected into pipe-work and will be detected emerging where
the leak is.
o Nuclear fission (power & atomic bomb)
 Uranium nuclei can be bombarded with neutrons and converted to other nuclei.
 The process occurs with the loss of massive amounts of energy in a chain reaction
 These huge amounts of energy can be used in a constructive or destructive way.
MOLECULES and IONS:
 Molecules are formed when a definite number of atoms are joined together by chemical bonds (covalent –
sharing of electrons).
 A molecule can consist of the atoms of only one element, or the atoms of many different elements – but
always in a fixed proportion.
o This means that molecules can be elements or compounds
o Molecules are usually formed between nonmetal elements
o Formula show the number of each type of atom present written as subscripts (lack of a subscript
means that only one type of that atom is present)
Examples of molecules:
Substance & Formula:
Hydrogen (H2), Oxygen (O2), Nitrogen (N2), Fluorine
(F2), Chlorine (Cl2), Iodine (I2), Bromine (Br2)
Water (H2O)
Ammonia (NH3)

Element or Compound:
Elements
Description:
Diatomic
Compound
Polyatomic
Compound
Polyatomic
Ions are formed when atoms gain or lose electrons causing the proton:electron ratio to become unbalanced
and the particle to become charged. These charged particles are termed ions.
o Positive ions (where the number of protons is greater than the number of electrons) are termed
cations
o Negative ions (where the number of electrons are greater than the number of protons) are termed
anions
o Metals tend to form cations and nonmetals tend to form anions
4
Name:
Date:
o
o
o
o
These oppositely charged ions form compounds by attracting one another (electrostatic attractive
forces – like charges repel and opposite charges attract)
An ion made up of only one type of atom is called a monatomic ion & an ion made up from more
than one type of atom is called a polyatomic ion.
 The bonds within polyatomic ions are covalent
Many atoms gain or lose enough electrons to have the same number of electrons as the nearest
noble gas (group 18)
The number of electrons an atom loses is relative to its position on the periodic table
Examples of ions:
Substance Formula/Symbol:
Sodium ion (Na+)
Cation or Anion:
Cation
Description:
Monatomic
Chloride ion (Cl-)
Anion
Monatomic
Carbonate ion (CO32-)
Anion
Polyatomic
Ammonium ion (NH4+)
Cation
Polyatomic
NOMENCLATURE (INORGANIC):
 Binary compounds are those formed between two elements. Inorganic binary compounds can have ionic
bonds and/or covalent bonds
o IONIC compounds consist of a metal and a nonmetal ionically bonded (ionic – transfer of
electrons)
 To determine the formula of an ionic compound the positive and the negative charges
must be balanced (no net charge)
 To name a binary compound of a metal and a nonmetal, the unmodified name of the
positive ion (cation – metal) is written first, followed by the root of the negative ion
(anion – nonmetal) with the ending modified to –ide
 If a transition metal is involved in the ionic compound, roman numerals are used to
indicate the oxidation number (example: iron (III) )
 ALSO – for metals with only two ions – the ion with the higher charge ends in –ic and
the ion with the lower charge ends in –ous
 Example – ferric & ferrous (aka: iron)

o
MUST KNOW FORMULAS & CHARGES OF COMMON IONS
MOLECULAR compounds consist of two nonmetals bonded covalently (covalent – sharing of
electrons).
 To name a molecular compound of two nonmetals, the unmodified name of the first
element is followed by the root of the second element with the ending modified to –ide.
In order to distinguish between several different compounds with the same elements
present use the following prefixes:
 Most common prefixes:
1. mono2. di3. tri4. tetra5. penta6. hexa7. hepta8. octa9. nona10. deca Examples: SO2, sulfur dioxide; BCl3, boron trichloride; CO, carbon monoxide
 Note that the prefix mono- is only applied to the second element present in such
compounds
5
Name:
Date:




Also, if the prefix ends with a or o and the element name begins with a or o then
the final vowel of the prefix is omitted
 Finally, some compounds have trivial names that have come to supercede their
systemic names (example: water not dihydrogen monoxide)
o BINARY ACIDS are formed when hydrogen ions combine with monatomic anions.
 For the purposes of nomenclature (naming) an acid can be defined as a compound that
produces hydrogen ions when it is dissolved in water.
 To name a binary acid, use the prefix hydro- followed by the other nonmetal name
modified to an –ic ending (example: HCl – hydrochloric acid)
Polyatomic ions are those where more than one element are combined together to create a species with a
charge. Some of these ions are names systemically, other names must be learned
o Polyatomic anions where oxygen is combined with another nonmetal are called oxyanions and
can be named systemically.
 In these oxyanions certain nonmetals (such as Cl, N, P, and S) form a series of oxyanions
containing different numbers of oxygen atoms.
 Their names are related to the number of oxygen atoms present.
 The ion with more oxygen atoms ends in –ate
 The ion with less oxygen atoms ends in –ite
 The ion with the most oxygen atoms begins with per- and ends in –ate
 The ion with the least oxygen ions begins with hypo- and ends in –ite
o EXAMPLE: ClO4- is perchlorate; ClO3- is chlorate; ClO2- is chlorite;
and ClO- is hypochlorite
 Some oxyanions contain hydrogen and are named accordingly (example: HPO 42-,
hydrogen phosphate)
 Also, the prefix thio- means that a sulfur atom has replaced an atom of oxygen in an
anion
 To name an ionic compound containing a polyatomic ion, the unmodified name of the
positive ion is written first followed by the unmodified name of the negative ion.
 Examples: NH4NO3, ammonium nitrate; K2CO3, potassium carbonate
Oxyacids are formed when hydrogen ions combine with polyatomic oxyanions giving a combination of
hydrogen, oxygen, and another nonmetal.
o To name an oxyacid use the name of the oxyanion and relace the –ite with –ous or the –ate with –
ic, then add the word acid. (example H2SO4, sulfuric acid)
Formula & Name of Oxyacid:
HClO
Hypochlorous acid
Formula & Name of Oxyanion:
ClOHypochlorite
HClO2
Chlorous acid
ClO2-
Chlorite
HClO3
Chloric acid
ClO3-
Chlorate
HClO4
Perchloric acid
ClO4-
Perchlorate
Hydrates are ionic formula units (compounds) with water associated with them.
o The water molecules are incorporated into the solid structure of the ions.
o To name a hydrate, use the normal name of the ionic compound followed by the term hydrate with
an appropriate prefix to show the number of water molecules per formula unit
 Example: CuSO4∙5H2O, copper (II) sulfate pentahydrate
o Strong heating can generally drive off the water in these salts. Once the water has been removed
the salts are said to be anhydrous (without water).
NOMENCLATURE (ORGANIC):
 Organic compounds are named according to IUPAC rules
o Simplest organic compounds are hydrocarbons
6
Name:
Date:

Four major classes of hydrocarbons: alkanes, alkenes, alkynes, aromatics
 Alkanes contain only single bonds
o Termed saturated
o General formula is CnH2n+2
 Alkenes contain at least one carbon-carbon double bonds
o Termed unsaturated & also referred to as olefins
o General formula is CnH2n
 Alkynes contain a carbon-carbon triple bond
o Also unsaturated
o General formula is CnH2n-2
 Aromatics have carbon atoms connected in a planer ring structure
o The carbons are linked by sigma & pi bonds
o Best known example is benzene (C6H6)

o
ALKANES:
 Straight-chain hydrocarbons
o Carbon atoms are joined in a continuous chain with no carbon atom
attached to more than two other carbon atoms
o Straight-chain hydrocarbons are not linear.
 Each carbon atom is tetrahedral, so the chains are bent
 Branched-chain hydrocarbons
o Possible for alkanes with four or more carbon atoms
o Structures with different branches can be written for the same formula
 These compounds are structural isomers having somewhat
different physical and chemical properties.
 Alkanes are very unreactive due to the strength of the C-C and C-H bonds.
 At room temperature alkanes do not react with acids, bases, or strong oxidizing
agents.
 Alkanes do undergo combustion in air (making them good fuels)
TO NAME ALKANES:
 Naming straight-chain alkanes varies according to the number of carbon atoms present in
the chain
 The names end in –ane
7
Name:
Date:

Prefix assigned denotes the number of carbon atoms

o
o
o
o
o
o
Find the longest chain and use it as the base name of the compound
 Groups attached to the main chain are termed substituents
 Number the carbon atoms in the longest chain starting with the end closest to a
substituent
 The preferred numbering will give substituents the lowest numbers
 Name and give the location of each substituent
 A substituent group formes by removing an H atom from an alkane is called an
alkyl group
 Alkyl groups are names by replacing the –ane ending with –yl
o Example: CH4 is methane, and a –CH3 group is a methyl group
 When two or more substituents are present, list them in alphabetical order
 When there are two or more of the same substituent, the number of that type of
substituent is indicated by a prefix (example: dimethyl indicates two methyl
group substituents)
Naming for CYCLOALKANES follows the same rules as alkanes except that the root name is
precede by the prefix cycloTO NAME ALKENES:
 Named the same way as alkanes with the suffix –ene replacing the –ane used for alkanes
 Location of double bond is indicated by a number
 If a substance has more than one double bond, the number of double bonds is denoted
with a prefix
 NOTE: only for alkenes containing more than 3 carbon atoms
TO NAME ALKYNES:
 Named in the same way as alkenes with the suffix –yne replacing the –ene used for
alkenes
Naming for CYCLOALKENES/-YNES:
 Number through the multiple bond towards the substituent
TO NAME AROMATICS:
 Special group of cyclic unsaturated hydrocarbons
 The delocalized pi electrons are usually represented by a circle in the center of the ring.
 Naming follows system for saturated ring systems
 Some aromatics can be substituents themselves in which case their naming is changed
 Example: simplest aromatic is benzene (C6H6). When benzene is a substituent it
is termed a phenyl group.
FUNCTIONAL GROUPS:
 A functional group is an atom, or groups of atoms, that occur together and as a whole
possess their own characteristic properties.
 Functional groups are the sites of reactivity in an organic molecule and they determine
the chemistry of a molecule.
8
Name:
Date:


The existence of each functional group is shown by adding a suffix and when
appropriate, a prefix to the name of the longest chain.
Common Functional Groups:
9