Download Atoms and the Periodic Table

Unit 3 – Lessons 1-4 “ATOMS and the PERIODIC TABLE” (Pages 153-208)
I. LESSON 1
A. The Atom
1. The basic UNIT of all MATTER is the ATOM
2. ATOM  is the SMALLEST particle of an ELEMENT retaining all the chemical PROPERTIES
of that ELEMENT
3. Democritus’s GREEK term “ATOMOS,” which means “not able to be divided,” is the ORIGIN of
the word, ATOM
*4. Individual ATOMS are so SMALL they can only be seen using a scanning tunneling
MICROSCOPE (STM), a special type of ELECTRON microscope
*a. Due to their SMALL size, huge numbers of ATOMS are used in the composition of very small
substances
*b. The average DIAMETER (distance across) of an ATOM is about 0.000000002 m or
2 x 10-9 m (scientific notation) or [2 nm (nanometers)]
*5. Scientific notation  is a system of writing NUMBERS with a LARGE number of place value
positions containing zeroes
*a. SCIENTIFIC NOTATION is a NUMBER written as a product with two FACTORS
1. The 1st FACTOR is a single digit whole number greater than 0, but less than 10
2. The 2nd FACTOR is a POWER of 10 in exponential form [(e.g.) “105” ]
*b. Steps for writing a NUMBER in a SCIENTIFIC NOTATION format:
1. Move the decimal POINT until the first NUMBER is “1” or between “1-10” but not 10
a. The first factor may be a MIXED-DECIMAL (e.g. 1.8 or 3.22)
2. Count the number of PLACE VALUE positions moved and express that total as an
EXPONENT
a. When you move the DECIMAL point to the left, the EXPONENT will be
expressed as a POSITIVE number
b. When you move the DECIMAL point to the RIGHT, the EXPONENT will
be expressed as a NEGATIVE number
3.(e.g.) Speed of light  30, 000,000,000 cm/sec = 3 x 10 10 cm/sec
Mass of a dust particle  0.000000000753 kg = 7.53 x 10 -10 kg
*c. A neutron’s mass is 168 hundred-octillionths of a kg
*1. (e.g.) Standard format: 0.00000000000000000000000000168 kg
Scientific notation: 1.68 x 10 -27
*d. An electron’s mass is 9,109,822 hundred-undecillionths of a kg
*1. (e.g.) Standard format: 0.0000000000000000000000000000009109822 kg
Scientific notation: 9.109822 x 10 -31
*e. The Earth’s circumference is 40 million meters
*1. (e.g.) Scientific notation: 40,000,000 = 4 x 107
f. A penny contains about “20” sextillion ATOMS of COPPER [Cu] and ZINC [Zn]
1. (e.g.) Scientific notation: 20,000,000,000,000,000,000,000 = 2 x 1022
B. Atomic Theory
1. Dalton’s THEORY of the ATOM having certain characteristics, form the BASIS of our understanding
of ATOMS today
a. All MATTER is composed of ATOMS
b. Atoms can NOT be CREATED, DIVIDED or DESTROYED
c. All the ATOMS of a specific ELEMENT are IDENTICAL
1. Explains why an ELEMENT always has the same PROPERTIES
d. Atoms of two or more ELEMENTS can COMBINE to form NEW substances
*e. Atoms of each ELEMENT have a unique MASS
*1. The number of PROTONS plus NEUTRONS in the NUCLEUS give an atom its
characteristic atomic MASS
*f. The MASSES of the ELEMENTS in a COMPOUND are ALWAYS in the same RATIO
*1. Since compounds have a specific chemical FORMULA, the ELEMENT(S) and the
number (or RATIO) of ATOMS per ELEMENT are always the SAME. Therefore,
the RATIO of the MASS of each ELEMENT is also always the SAME
*a. (e.g.) calcium oxide [Ca5O2] = 5 (Ca amu) + 2 (O amu) = 232 amu
5 (40 amu) + 2 (16 amu)
(200 amu) + (32 amu)
2. Thomson provided evidence of ATOMS having NEGATIVELY charged particles = ELECTRONS
3. Rutherford’s experiments provided evidence that the ATOM contained a DENSE, central
NUCLEUS composed of POSITIVELY charged, SUBATOMIC particles called PROTONS
4. Chadwick discovered that the ATOM contained UNCHARGED particles called NEUTRONS
5. Bohr’s observations provided ATOM behavior based on ELECTRON movement around the
NUCLEUS on circular ORBITAL paths (ENERGY levels)
6. Today:
a. ELECTRON CLOUDS  the area around the NUCLEUS where ELECTRONS move
*b. QUARKS  the sub-subatomic particles that make up PROTONS and NEUTRONS
C. Parts of an Atom
1. Within the NUCLEUS:
a. NUCLEUS:
 refers to the dense, inner most CORE of an ATOM containing the PROTONS
(POSITIVELY charged particles) and the NEUTRONS (particles that have NO CHARGE)
*1. the NUCLEUS is the most MASSIVE part of an ATOM
b. PROTON:
 POSITIVE electrical CHARGE
*1. written as: one proton = 1+; atom with “14” PROTONS = 14+

this MASSIVE subatomic particle along with the NEUTRON count gives an ATOM its
atomic MASS number (measured in “amu” units or UNIFIED atomic mass units “u”)
*1. mass of a PROTON = 0.0000000000000000000000000017 u = 17 ten-octillionths
u = 1.7 x 10 -27 u (in SCIENTIFIC NOTATION)

number of PROTONS (atomic number) determines the IDENTITY of the ELEMENT
*1. ATOMIC number  is the PROTON count; unique for each ELEMENT
*2. (e.g.) atomic number = 5; PROTON count = 5; element = BORON [B]

proton COUNT can NOT change for an ELEMENT or you have a different ELEMENT
*1. (e.g.) PROTON count = 10; element = NEON [Ne]
“1” less proton than neon = 9; element = FLUORINE [F]
“1” more proton than neon = 11; element = SODIUM [Na]
c. NEUTRON:
 NO electrical CHARGE
*1. (e.g.) beryllium [Be]: 4 protons = “4+” and 5 NEUTRONS = “4+” NET charge

this MASSIVE subatomic particle along with the PROTON count gives an ATOM its
atomic MASS number (measured in “amu” units or UNIFIED atomic mass “u”)
*1. mass of a NEUTRON = slightly more than a PROTON
*2. (e.g.) beryllium [Be]: 4 protons = “4+” and 5 NEUTRONS = “9” u (or “9” amu)

neutron COUNT CAN change for an ELEMENT forming ISOTOPES
*1. Steps to finding the NEUTRON count from a PERIODIC table
*a. Find the atomic number (PROTON count)
*b. Find the average atomic mass number (mixed decimal on Periodic Table)
*c. ROUND the average atomic mass number to the nearest whole number
*d. SUBTRACT the atomic number from the atomic mass number
*e. (e.g.) manganese [Mn]  atomic number = 25
average atomic mass # = 54.94 ~ 55 (atomic mass #)
manganese [Mn] = “30” neutrons
*2. ISOTOPE  atoms of the same ELEMENT with the same PROTON count
(atomic number), but have a DIFFERENT number of NEUTRONS
and therefore have a DIFFERENT atomic MASS number
*a. hydrogen (protium) [H]  proton = 1+; neutron = 0; 1 amu
*b. hydrogen (deuterium) [H]  proton = 1+; neutron = 1; 2 amu (isotope)
*c. hydrogen (tritium) [H]  proton = 1+; neutron = 2; 3 amu (isotope)
2. Outside the NUCLEUS:
a. ELECTRON:
 NEGATIVE electrical CHARGE
*1. written as: one electron = 1-; atom with “14” ELECTRONS = “14-”
2. the charges of PROTONS and ELECTRONS are OPPOSITE but EQUAL, which
makes the ATOM electrically NEUTRAL (NET charge = “0”)
*3. (e.g.) beryllium [Be]: 4 protons = “4+” + “4-” ELECTRONS = “0” net charge

has relatively NO MASS; the ELECTRONS are NEVER used to determine the atomic
MASS of an ELEMENT
*1. MASS of an ELECTRON is 1,860 times LESS than a PROTON or NEUTRON
*2. (e.g.) lithium [Li]: protons = 3+; electrons = 3-; neutrons = 4 = 7 u (or 7 amu)

are in constant MOTION around the OUTSIDE of the nucleus within the electron CLOUD
*1. HEISENBERG RULE  states that it is NOT possible to determine the exact
LOCATION and SPEED of an ELECTRON simultaneously
*2. ELECTRON models show their movement on paths called energy ORBITALS (or
energy LEVELS or energy SHELLS)

electron COUNT CAN change for an ELEMENT when forming IONIC bonds by
transferring (giving up or taking on) ELECTRONS, producing IONS
*1. ION  an electrically charged ATOM (having a “+” or “-” charge) due to
ELECTRONS being TRANSFERRED (gained or lost)
*2. IONIC bond  the FORCE that attracts OPPOSITELY charged IONS and
CHEMICALLY holds them together
*a. sodium [Na] will lose “1” ELECTRON to become stable = “Na+”
*b. chlorine [Cl] will gain “1” ELECTRON to become stable = “Cl-”
*3. STABLE  the outer-most, VALENCE energy ORBITAL contains the
MAXIMUM number of valence ELECTRONS
*a. helium [He] = 1st energy ORBITAL contains “2” electrons (maximum)
*b. krypton [Kr] = 4th energy ORBITAL contains “8” electrons (maximum)
D. Atomic Number verses Atomic Mass Number
1. Different COMBINATIONS of PROTONS, NEUTRONS and ELECTRONS produce ATOMS
with different PROPERTIES (ELEMENTS)
a. These different ATOMS chemically COMBINE to form different, NEW substances
(COMPOUNDS)
2. The number of PROTONS distinguishes each ATOM from the other 118 ELEMENTS
3. ATOMIC NUMBER  the number of PROTONS in the NUCLEUS of an ATOM
a. (e.g.) iron [Fe] atomic number = 26; proton count = 26+
tin [Sn] atomic number = 50; proton count = 50+
mercury [Hg] (a LIQUID metal) atomic number = 80; proton count = 80+
bromine [Br] (only liquid NON-METAL) atomic number = 35; proton count = 35+
californium [Cf] (period 7; actinides) atomic number = 98; proton count = 98+
4. The ATOM of a specific ELEMENT always has the SAME number of PROTONS, but it does
NOT always have the same number of NEUTRONS
a. (e.g.) chlorine [Cl] ALWAYS has 17+ PROTONS, but it could have 18 or 20 NEUTRONS
b. chlorine-35 [Cl] with 17+ PROTONS and 18 NEUTRONS = ISOTOPE of CHLORINE
chlorine-37 [Cl] with 17+ PROTONS and 20 NEUTRONS = ISOTOPE of CHLORINE
5. Atomic MASS NUMBER  is TOTAL number of PROTONS plus NEUTRONS in the
NUCLEUS of a specific ATOM
a. (e.g.) silver [Ag] (period 5; group 11) proton count = 47+; neutron count = 61 = 108 u
argon [Ar] (period 3; group 18) proton count = 18+; neutron count = 22 = 40 u
arsenic [As] (period 4; group 15) proton count = 33+; neutron count = 42 = 75 u
lead [Pb] (period 6; group 14) proton count = 82+; neutron count = 125 = 207 u
uranium [U] (period 7; actinides) proton count = 92+; neutron count = 146 = 238 u
*6. AVERAGE ATOMIC MASS  is the weighted average of the MASSES of all the naturally
occurring ISOTOPES of an ELEMENT
*a. generally MIXED decimals on the Periodic Table
*1. Exception: RADIOACTIVE elements’ most common ISOTOPES (whole numbers)
b. (e.g.) potassium [K] (period 4; group 1) average atomic mass = 39.10
phosphorus [P] (period 3; group 15) average atomic mass = 30.97
gold [Au] (period 6; group 11) average atomic mass = 196.97
praseodymium [Pr] (period 6; lanthanides) average atomic mass = 140.91
mendelevium [Md] (period 7; actinides) average atomic mass = (258)
II. LESSON 2
A. The History of the Periodic Table
*1. In 1828, Dobereiner made one of the earliest attempts to “list” the ELEMENTS; proposed the Law
of Triads  stated there were groups of “3” ELEMENTS where the middle ELEMENT’S atomic
MASS was the average of the other “2” ELEMENTS
*a. (e.g.) calcium [Ca] = 40.08; strontium [Sr] = 87.62; barium [Ba] = 137.33
*2. In 1856, Newlands proposed the Law of Octaves  stated that some of the 56 ELEMENTS whose
atomic MASS differed by some multiple of EIGHT had similar PROPERTIES
*a. this law supported the “Octet rule” for valence ELECTRONS and chemical BONDING
3. The first PERIODIC TABLE was created and published by MENDELEEV in 1869, who arranged the
63 existing ELEMENTS in order by their increasing ATOMIC MASS number
a. “PERIODIC”  means having a regular, REPEATING pattern and it means a “listing”
b. Mendeleev used the “periodic patterns” to PREDICT future ELEMENTS
4. In 1914, Moseley re-organized the PERIODIC Table according to each element’s increasing
ATOMIC number (PROTON count) from left to RIGHT rows (PERIODS), rather than by their
atomic mass number
5. Today the PERIODIC TABLE is a valuable tool showing many PATTERNS among the 118
elements’ PROPERTIES
*a. Some ELEMENTS are very reactive and form COMPOUNDS easily, while others are less
REACTIVE, and still others do NOT form COMPOUNDS at all
B. Information on the Periodic Table
1. Each SQUARE on the PERIODIC TABLE represents an ELEMENT
2. Each SQUARE generally includes the element’s:
 ATOMIC number  corresponds to the number of PROTONS for that element
*a. all ISOTOPES (different NEUTRON count) of that ELEMENT, but have the SAME
ATOMIC number
*b. all ELEMENTS on the periodic table are NEUTRAL (same number of PROTONS
and ELECTRONS), the ATOMIC number also gives the ELECTRON count
 CHEMICAL symbol  an ABBREVIATED form of the element’s NAME
*a. 1st letter is always CAPITALIZED any other letter(s) are ALWAYS lower-case
*b. some CHEMICAL symbols are taken directly from the ELEMENT name, others are
based on their their LATIN, GREEK or Arabic names
*c. some CHEMICAL symbols have “3” letters to represent their TEMPORARY name
*1. (e.g.) yttrium [Y]; francium [Fr]; tungsten [W] (wolfram); Ununtrium [Uut]
 CHEMICAL name  the name of the ELEMENT
*a. sources for names come from SCIENTISTS; LOCATIONS; UNIVERSITIES; etc…
*1. (e.g.) einsteinium [Es]; europium [Eu]; berkelium [Bk]
 AVERAGE ATOMIC MASS  the weighted AVERAGE of the MASSES of all the
natural occurring ISOTOPES of that ELEMENT
*a. the unit label for ATOMIC MASS is “u” or “amu”
*b. by rounding the AVERAGE ATOMIC MASS and subtracting the ATOMIC
NUMBER from it, the NEUTRON count can be calculated for an ELEMENT
*c. The ATOMIC MASS number of an ELEMENT can vary, because the ATOM of an
ELEMENT can have varying numbers of NEUTRONS (ISOTOPES)
*1. (e.g.) Carbon ATOMS must always contain 6 PROTONS, but they may contain
anywhere from 5-8 NEUTRONS
*d. ISOTOPES  are ATOMS with the SAME number of PROTONS, but with
DIFFERENT numbers of NEUTRONS.
*1. (e.g.) Carbon-12  6 PROTONS and 6 NEUTRONS
*a. Format:
Atomic Mass Number  12 u
Atomic Number  6
Carbon-14  6 protons and 8 NEUTRONS
*a. Format:
Atomic Mass Number  14 u
Atomic Number  6
3. Inside the Square
a. Xenon [Xe] square on the PERIODIC TABLE:
54
Xe
xenon
131.29
ATOMIC NUMBER
CHEMICAL SYMBOL
CHEMICAL NAME
AVERAGE ATOMIC MASS
*C. Organization of the Periodic Table
*1. The ELEMENTS are arranged in order of INCREASING ATOMIC numbers from LEFT to
RIGHT and from TOP to BOTTOM
*2. The ELEMENTS are arranged in “7” horizontal rows called PERIODS and “18” vertical columns
called GROUPS (or FAMILIES)
*a. Extensions of PERIOD 6 (lanthanides) & PERIOD 7 (actinides) are below the periodic table
*3. An element’s PROPERTIES can be predicted based on its LOCATION on the periodic table
*a. to the right, left or bordering the ZIGZAG line indicates the “classification” of the
ELEMENT as a METAL, NON-METAL or METALLOID
*b. the element’s PERIOD number and GROUP number indicates what “properties” the
ELEMENT will have
*4. The PERIODIC TABLE is “coded” (colors and symbols) to indicate the ELEMENTS’ STATE of
matter (solid, liquid or gas) at ROOM temperature (250 C) and to assist in identifying its
CLASSIFICATION (metal, non-metal or METALLOID)
D. Metals, Non-metals and Metalloids
1. The “3” major CLASSIFICATION categories on the PERIODIC TABLE
2. The bolded, ZIGZAG line assists in the IDENTIFICATION as to what class the ELEMENT belongs
a. METALS are to the LEFT of the LINE (Exception: hydrogen [H])
1. PROPERTIES: LUSTER (shiny); CONDUCT electricity and HEAT; most are SOLID
at 250 C; MALLEABLE (easily formed into different shapes) and DUCTILE (able to be
made into a WIRE)
b. NON-METALS are to the RIGHT of the ZIGZAG line
1. PROPERTIES: OPPOSITE of METALS (dull, POOR conductors of HEAT and
electricity, BRITTLE)
c. METALLOIDS border the ZIGZAG line
1. SIX metalloids: boron [B]; silicon [Si]; germanium [Ge]; arsenic [As]; antimony [Sb]
and tellurium [Te]
1. PROPERTIES: have properties of both METALS and NON-METALS
*a. (e.g.) Conduct HEAT: NOT as good as METALS; better than NON-METALS
E. The Groups
1. GROUPS  the “18” vertical COLUMNS forming the main body of the PERIODIC TABLE
2. The GROUPS have a FAMILY name, some based on the first ELEMENT in that COLUMN
*a. Group 1 = ALKALI metals; Group 2 = ALKALINE EARTH metals; Groups 3-12 =
TRANSITION metals (Group 11 = COINAGE metals); Group 13 = BORON family;
Group 14 = CARBON family; Group 15 = NITROGEN family; Group 16 = OXYGEN
family; Group 17 = HALOGENS; Group 18 = NOBLE gases (or inert gases)
3. The ELEMENTS in each GROUP have similar PHYSICAL and CHEMICAL properties because
they have the SAME number of VALENCE electrons
*a. VALENCE ELECTRON(S)  are the ELECTRONS found on the OUTER-MOST portion
of the electron CLOUD (energy orbital FARTHEST from the NUCLEUS) of an ATOM,
which allows them to participate in chemical BONDING since they are farthest away from
the attractive FORCE of the NUCLEUS
*1. Group 1 are all METALS that react VIOLENTLY with WATER
*a. From (Li – Fr) the alkali metals are ONLY found in COMPOUNDS
*b. Group “1” valence electron count = “1”
*c. Hydrogen [H] – which is NOT one of the alkali metals due to its different
chemical PROPERTIES, is the simplest ELEMENT
*1. 90% of the ATOMS of the universe are [H], but only 1% of the MASS
of the Earth’s crust, OCEANS, and ATMOSPHERE is made up of [H]
*2. Group 2, the alkaline earth metals, are good CONDUCTORS of electricity
*a. From (Be – Ra) the alkaline earth metals = NOT as REACTIVE as Group 1
*b. Group “2” valence electron count = “2”
*3. Groups 3-12, the transition metals, bridge the REACTIVE metals with the LESS
reactive metals; good conductors of HEAT and ELECTRICITY; contain the most
FAMILIAR metals; have LUSTER (shiny); and contains the iron [Fe] an essential
METAL for the human BODY to make hemoglobin, which is necessary for carrying
OXYGEN [O] in the bloodstream
*a. transition metals do NOT follow the “Octet Rule” of needing “8” VALENCE
electrons to be STABLE
*b. STABLE  an ATOM with the MAXIMUM number of ELECTRONS in its
VALENCE orbital
*c. Group 11, also known as COINAGE metals (copper [Cu], silver [Ag], gold
[Au]), are slow to REACT with water
*4. Groups 13-16 include METALS, NON-METALS, and the METALLOIDS and are
composed of the BORON family (Group 13 w/ “3” valence electrons); CARBON
family (Group 14 w/ “4” valence electrons); the NITROGEN family (Group 15 w/ “5”
valence electrons) and the OXYGEN family (Group 16 w/ “6” valence electrons)
*5. Group 17, the halogen family, reacts VIOLENTLY with the elements in GROUP 1,
and therefore in their pure form are DANGEROUS to HUMANS, while in compounds
they are very useful [(e.g.) Individually “Na” (explosive metal) and “Cl” (poisonous
gas), but as a COMPOUND it forms TABLE SALT [NaCl]
*a. Group “17” valence electron count = “7”
*6. Group 18, the noble or inert gases, rarely REACT at all, since they are very stable
ELEMENTS due to their VALENCE electron orbitals being FULL
*a. Group “18” valence electron count = “8”; Exception: helium [He] = “2”
F. The Periods
1. PERIOD  is one of the 7 horizontal ROWS going across the PERIODIC TABLE containing a
series of DIFFERENT types of ELEMENTS from different GROUPS (or families)
a. The ELEMENTS in each PERIOD gradually change their physical and CHEMICAL
PROPERTIES as you move from LEFT to RIGHT in PREDICTABLE ways:
1. (e.g.) atomic SIZE decreases (moving LEFT to RIGHT)
2. (e.g.) DENSITY of ELEMENTS is the LEAST dense on the LEFT and RIGHT sides
of a PERIOD, while being MOST dense in the MIDDLE of the PERIOD
*a. osmium [Os] (period 6; group8) has the HIGHEST known DENSITY
*b. Period 1 contains two ELEMENTS: [H] and [He] and is referred to as a “short period”
Periods 2 & 3 each contain 8 elements and are also “short periods”
Periods 4 & 5 each contain 18 elements.
Periods 6 & 7 each contain 32 ELEMENTS and therefore both have a portion which is
SEPARATED from the main body of the PERIODIC TABLE
*1. Period 6’s ELEMENTS (58-71) are the LANTHANIDES, which are MALLEABLE
(easily shaped) and used to make ALLOYS (mixture of metals; metals & non-metals)
*2. Period 7’s ELEMENTS (90-103) are the ACTINIDES, many are ARTIFICALLY
made in labs, are unstable so they only last for a fraction of a SECOND after being
made, and exist in small amounts (Exception: thorium (Th) and uranium [U])
*3. The LANTHANIDES and ACTINIDES are also known as the “rare earth elements”
4. The LANTHANIDES and ACTINIDES allow for the MAIN BODY of the
PERIODIC TABLE to be NARROWER
5. The LANTHANIDES and ACTINIDES elements also are arranged in INCREASING
order by their ATOMIC number
*c. PERIOD number also indicates the number of energy ORBITALS an ATOM of a specific
ELEMENT will have and for ELECTRON configuration its PRINCIPAL QUANTUM
number
 Period 1 = hydrogen [H] and helium [He] only have “1” energy ORBITAL (type “s”)
and their Principal QUANTUM number = “1”
 Period 2 has 2 energy ORBITALS (types: “s” and “p”); Principal quantum # = “2”
 Period 3 has 3 energy ORBITALS (types: “s” and “p”); Principal quantum # = “3”
 Period 4 has 4 energy ORBITALS (types: s; p and “d”); Principal quantum # = “4”
 Period 5 has 5 energy ORBITALS (types: s; p and “d”); Principal quantum # = “5”
 Period 6 has 6 energy ORBITALS (types: s; p; d and “f”); Principal quantum # = “6”
 Period 7 has 7 energy ORBITALS (types: s; p; d and “f”); Principal quantum # = “7”
III. LESSON 3
A. Chemical Bonding and Chemical Changes
1. There are “118” elements and about “92” of these ELEMENTS are known to exist in NATURE and
form all MATTER due to CHEMICAL BONDING
2. CHEMICAL BOND  is a FORCE (or interaction) that holds ATOMS or IONS together forming
MOLECULES
3. MOLECULE  is the SMALLEST unit of a COMPOUND where two of more ATOMS are
CHEMICALLY joined together by chemical BONDS
a. some MOLECULES are extremely SMALL [(e.g.) water molecule [H2O] = “3” ATOMS]
b. some are LARGE [(e.g.) DNA molecule [deoxyribonucleic acid] = “BILLIONS” of atoms]
c. Due to the chemical BONDING of joining ATOMS together into a MOLECULE, the
molecule acts as a single UNIT
4. CHEMICAL BONDING always involves chemical CHANGES by REARRANGING the ORDER
of the ATOMS either by JOINING the atoms together or by BREAKING them apart
*a. (e.g.) 2 H2 + O2  2 H2O (SYNTHESIS chemical reaction)
“2” diatomic MOLECULES of HYDROGEN gas COMBINE with “1” DIATOMIC
MOLECULE of OXYGEN gas, which YIELDS “2” MOLECULES of WATER
*b. (e.g.) 2 H2O  2 H2 + O2 (DECOMPOSITION chemical reaction)
“2” MOLECULES of WATER (decompose), which YIELDS “2” DIATOMIC
MOLECULES of HYDROGEN gas AND “1” diatomic MOLECULE of oxygen GAS
*5. When substances undergo a CHEMICAL change, NEW substances with different PHYSICAL and
chemical PROPERTIES are formed
*a. (e.g.) The properties of the REACTANTS are completely DIFFERENT from those of the
PRODUCT, yet hydrogen [H] and oxygen [O] are the only ELEMENTS used in the
formation of the 2 water MOLECULES
2 H2 + O2 (the REACTANTS)  2 H2O (the product)
*b. The CHEMICAL change did NOT create or DESTROY any ATOMS (Law of
CONSERVATION of MATTER), it simply caused the hydrogen [H] and oxygen [O] ATOMS
to be REARRANGED
Balancing the equation: 2 H2 + O2 (the REACTANTS)  2 H2O (the PRODUCT)
Reactants (original substances):
Elements:
Number of atoms/element:
H
=
4
O
=
2

Elements:
H
O
Products (new substances):
Number of atoms/element:
=
4
=
2
B. Atomic Models
*1. Using Models
*a. MODEL  is a representation of how something LOOKS and/or WORKS
*b. The PARTICLE model of MATTER is used to study ATOMS and MOLECULES
*1. The study of ATOMS explains the various PROPERTIES of ELEMENTS
*2. The study of ELEMENTS explains the various PROPERTIES of COMPOUNDS
2. MODELS of ATOMS are used to show how atoms BEHAVE and how ELECTRONS are involved
in chemical BONDING
3. There are “3” different MODELS used to represent ATOMS: ELECTRON CLOUD model, the
BOHR model and the SPACE-FILLED model
a. ELECTRON CLOUD model shows the atomic NUCLEUS in the CENTER of the atom and
the indistinct, CLOUD-like region around the NUCLEUS
1. ELECTRONS move through this CLOUD-like region
*2. Model re-enforces the HEISENBERG rule that the POSITION and trajectory of
moving ELECTRONS can NOT be precisely determined at the SAME instant
3. ELECTRON CLOUD models do NOT show the NUMBER of ELECTRONS
b. BOHR model shows the number of ENERGY levels and the number of ELECTRONS per
ENERGY level (or energy ORBITAL)
1. ELECTRONS are represented by DOTS
2. ENERGY levels are represented by RINGS
3. The BOHR model does NOT show the TRUE location of the ELECTRONS, but it
helps explain the chemical BONDING of ATOMS and their chemical PROPERTIES
c. SPACE-FILLED model represent ATOMS as SOLID SPHERES to show how they are
JOINED together in substances
1. The SPACE-FILLED model does NOT show the PARTS that make up the ATOMS
C. Valence Electrons and Chemical Bonds
*1. VALENCE electrons explain BONDING power because they can either be SHARED or
TRANSFERRED (lost/GAINED) between other ATOMS
*2. BONDING power  refers to the number of chemical BONDS an ELEMENT can form during a
CHEMICAL change
3. VALENCE ELECTRON  are the ELECTRONS that occupy the atom’s OUTER-MOST
energy level (valence electron ORBITAL) and have the LEAST
amount of ATTRACTIVE force TO the NUCLEUS
*a. They are the ONLY electrons that can be SHARED or TRANSFERRED (lost/GAINED)
*b. The AMOUNT of VALENCE electrons in the VALENCE electron orbital determines
whether the ATOM of an ELEMENT will give UP, take ON or SHARE electrons
*c. The number of VALENCE electrons an element has INCREASES from LEFT to RIGHT
across a PERIOD
4. VALENCE ELECTRON ORBITAL  is the atom’s OUTER-MOST energy orbital that is
FARTHEST from the ATTRACTIVE force TO the NUCLEUS
*5. Guidelines to follow for the amount of ELECTRONS placed on energy ORBITALS are:
*a. 1st energy orbital = “2” electrons (maximum); 2nd energy orbital = “8” electrons
(maximum); 3rd energy orbital = “18” electrons (maximum); 4th energy orbital = “32”
electrons (maximum); 5th energy orbital = “64” electrons (maximum); etc…
*b. The “18-Electron Rule”: 1st energy orbital = “2” electrons (max); 2nd energy orbital =
“8” electrons (maximum); 3rd energy orbital = “18” electrons (maximum); all other energy
orbitals “max” out at “18” electrons and are then considered STABLE
*c. The “Octet Rule” (or LEWIS RULE of Eight): states that ATOMS combine to form
MOLECULES by losing, GAINING or sharing VALENCE electrons until they attain “8”
VALENCE electrons and are considered STABLE: 1st energy orbital = “2” electrons
(maximum); 2nd energy orbital = “8” electrons (maximum); all other energy orbitals
“max” out at “8” electrons
6. The PERIODIC table (using the OCTET Rule) can also indicate the number of VALENCE
electrons for an ELEMENT
a. The ELEMENTS in each GROUP of the periodic table have the SAME NUMBER and
arrangement of VALENCE ELECTRONS in their VALENCE electron orbital
b. The GROUP number in GROUPS 1-2 and 13-18 can be used to find the NUMBER of
VALENCE electrons in their VALENCE electron orbital
1. EXCEPTIONS:
a. the transition metals (GROUPS 3-12) have either 1 or 2 VALENCE electrons
b. helium [He] is in GROUP 18, but only has “2” VALENCE electrons due to its
single VALENCE ELECTRON ORBITAL
2. Group 1 (ALKALI metals) - all the ELEMENTS have “1” valence electron
[H = SHARES; Li-Fr = TRANSFER by GIVING it up]
Group 2 (alkaline earth metals) - all the ELEMENTS have “2” valence electrons
[TRANSFER by GIVING them up]
Groups 3-12 (transition metals) - most the ELEMENTS have 1 or 2 valence electrons
[TRANSFER by GIVING them up]
Group 13 (BORON family) - all the ELEMENTS have “3” valence electrons
[TRANSFER by GIVING them up]
Group 14 (CARBON family) - all the ELEMENTS have “4” valence electrons
[SHARE; especially CARBON]
Group 15 (NITROGEN family) - all the ELEMENTS have “5” valence electrons
[TRANSFER by TAKING on electrons given up by other atoms]
Group 16 (OXYGEN family) - all the ELEMENTS have “6” valence electrons
[TRANSFER by TAKING on electrons given up by other atoms]
Group 17 (HALOGEN family) - all the ELEMENTS have “7” valence electrons
[TRANSFER by TAKING on electrons given up by other atoms]
Group 18 (NOBLE or INERT gases) - all the ELEMENTS have 8 valence electrons
(EXCEPTION: HELIUM [He] has “2” valence electrons)
[All STABLE elements; do NOT TRANSFER or SHARE their valence electrons]
7. ATOMS want to be STABLE (valence orbital is FULL) but if their VALENCE orbital does NOT
contain the maximum number of VALENCE electrons they will form chemical BONDS with other
ATOMS by TRANSFERRING (gain/lose) or SHARING VALENCE electrons
a. Forming BONDS allows ATOMS to FILL their VALENCE orbitals
*b. (e.g.) Group 1 (ALKALI metals) having “1” VALENCE electrons (very REACTIVE
elements) want to BOND with elements from Group 17 (HALOGEN family) that have
“7” VALENCE electrons (LESS reactive elements)
c. (e.g.) sodium [Na] (Group 1) has “1” VALENCE electron (LOSE “1” electron) = Na+
chlorine [Cl] (Group 17) has “7” VALENCE electrons (GAIN “1” electron) = ClCHEMICAL bond (ionic bond) by TRANSFERRING = NaCl (IONIC compound)
*D. Valence Electrons and Electron Configuration
*1. ELECTRON CONFIGURATION  a term for how ELECTRONS are arranged in an ATOM
*2. ELECTRON movement around the NUCLEUS is NOT in circular paths (Bohr’s model), but in
CLOUD-like zones (ORBITALS)
*3. HEISENBERG Uncertainty Rule  states that it is IMPOSSIBLE to know the LOCATION and
SPEED of the ELECTRONS simultaneously
*4. “ORBITALS” are where the ELECTRONS are LIKELY to be FOUND
*a. ORBITALS are GROUPED and IDENTIFIED according to their SHAPES: s, p, d and f
(also known as SUBSHELLS)
 “s”  “sharp”
 “p”  “principal”
 “d”  “diffuse”
 “f”  “fundamental”
*5. QUANTUM numbers  the “4” numbers that describe the atomic ORBITALS: n, l, ml, ms
*a. PRINCIPAL QUANTUM number (n)  is the WHOLE number which tells the overall
ENERGY of the ELECTRON and its LOCATION
 “n” is located in front of the PERIOD number and is the SAME as the PERIOD #
 the HIGHER the PRINCIPAL QUANTUM number the MORE electron energy
 FARTHER away from the NUCLEUS the MORE electron energy
 ELECTRONS with more ENERGY can occupy different types of ORBITALS
 LOWER energy ORBITALS fill FIRST (Exception: TRANSITION metals)
*b. QUANTUM number (l)  represents the orbital SHAPE and refers the “GROUPS”
(1-18) on the PERIODIC table
 Type “s” = spherical shape; “s” block = GROUPS “1 & 2” and HELIUM [He]
 Type “p” = dumbbell shape; “p” block = GROUPS “13 - 18” (without [He])
 Type “d” = many complicated shapes; “d” block = GROUPS “3 - 12”
 Type “f” = more complicated shapes; “f” block = (lanthanides & actinides)
*c. QUANTUM number (ml)  represents the orientation of the ORBITAL which only
effects the “p, d and f” ORBITALS
*d. QUANTUM number (ms)  represents the SPIN of the ELECTRON on its axis in one of
two possible directions
*1. An ORBITAL can only accommodate “2” ELECTRONS each since they are
SPINNING in an OPPOSITE direction
*2. Each ORBITAL can HOLD a MAXIMUM (based on “2” ELECTRONS per
ORBITAL  ms) of:
 Type “s” = has “1” ORBITALS = TOTAL “2” ELECTRONS
 Type “p” = has “3” ORBITALS = TOTAL “6” ELECTRONS
 Type “d” = has “5” ORBITALS = TOTAL “10” ELECTRONS
 Type “f” = has “7” ORBITALS = TOTAL “14” ELECTRONS
Orbital Type
[quantum # (l)]
s
Number of Orbitals
[(m2) max # of 2
electrons per subshell]
1
Maximum # of
Electrons per Type
of Orbital
2
p
3
6
d
5
10
f
7
14
*3. HUND RULE of MAXIMUM MULTIPLICITY (ms)  states that you MUST put
“1” ELECTRON in each ORBITAL of a subshell BEFORE you begin to
DOUBLE (or PAIR) the ELECTRONS up
*a. Each “p” ORBITAL (x, y and z) MUST receive “1” ELECTRON each
BEFORE electron “PAIRING” can occur
*b. (e.g.) oxygen [O]  ATOMIC number (PROTON count) = “8”
ELECTRON count (NEUTRAL atom) = “8”
Each ARROW represents “1” ELECTRON
1s
2s
2p (x)
2p (y)
2p (z)
*6. Writing FORMAT of ELECTRON CONFIGURATION for an ELEMENT: n, l, ml, ms
*a. 1st  write the PRINCIPAL QUANTUM number (“n”) which represents the energy
ORBITAL ( or “PERIOD” number)
nd
*b. 2  write the ORBITAL letter (“l”) representing the “ORBITAL type” (s, p, d or f)
*c. 3rd  write the “EXPONENT” which represents the number of “ELECTRONS” in that
ORBITAL
*d. The “ORDER” for filling ORBITALS: 1s, 2s, 2p, 3s, 3p, 4s…
For the order of filling orbitals, begin at the base of each
arrow starting with “1 s”, follow it all the way to the point
and begin at the base of the next arrow.
Order would be: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, etc…
n1
n2
n3
n4
n5
n6
n7
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
4f
5f
*e. (e.g.) Writing FORMAT of ELECTRON CONFIGURATION for:
*1. lithium [Li]  ATOMIC number (PROTON count) = “3”
ELECTRON count (NEUTRAL atom) = “3”
ELECTRON CONFIGURATION: ___1s2 2s1_______________
(drawing)
*2. carbon [C]  ATOMIC number (PROTON count) = “6”
ELECTRON count (NEUTRAL atom) = “6”
ELECTRON CONFIGURATION: ____1s2 2s2_2p2___________
(drawing)
*3. sodium [Na]  ATOMIC number (PROTON count) = “11”
ELECTRON count (NEUTRAL atom) = “11”
ELECTRON CONFIGURATION: ____1s2 2s2_2p6_3s1_______
(drawing)
*8. The VALENCE electrons in ELECTRON CONFIGURATION refer to ALL the electrons in the
HIGHEST PRINCIPAL QUANTUM number ORBITAL(S)
*a. (e.g.) nitrogen [N]  ATOMIC number (PROTON count) = “7”
ELECTRON count (NEUTRAL atom) = “7”
Electron Configuration: 1s2 2s 2 2p 3
Valence electron count: “5”
*b. (e.g.) sodium [Na]  ATOMIC number (PROTON count) = “11”
ELECTRON count (NEUTRAL atom) = “11”
Electron Configuration: 1s2 2s2 2p6 3s 1
Valence electron count: “1”
*c. For GROUPS 1, 2, 13-18 (EXCEPTION: helium [He]) the GROUP number will help to
determine the VALENCE ELECTRON count
*1. (e.g.) lithium [Li] = Group “1”  “1” valence electron
carbon [C] = Group “14”  “4” valence electron
neon [N] = Group “18”  “8” valence electron
magnesium [Mg] = Group “2”  “2” valence electron
helium [He] = Group “18”  “2” valence electron (exception)
*9. Information that can be determined by reading the ELECTRON CONFIGURATION of an
ELEMENT: ATOMIC number (or PROTON count); ELECTRON count (NEUTRAL atom);
PRINCIPAL QUANTUM (or PERIOD number); VALENCE electron count (or GROUP
number); NEUTRAL (same number of PROTONS as ELECTRONS); STABLE ( VALENCE
orbital is FULL); Type of BOND (IONIC, COVALENT or INERT)
*a. (e.g.) sulfur [S]  Electron Configuration: 1s2 2s2 2p6 3s2 3p4
 Atomic number/Proton or Electron count (sum of ALL the exponents) = “16”
Principal Quantum or Period number (highest LARGE number) = “3”
Valence Electron count (sum of exponents of highest LARGE number) = “6”
 Group number (sum of exponents of highest LARGE number) = “6”
 Neutral (when proton and electron count are the SAME number) = “YES”
 Stable (when valence electron count is “8” [“2” for helium]) = “NO”
 Type of Bond (ionic: transfer; covalent: share; inert: Grp18) = “IONIC”
IV. LESSON 4
A. Ionic Bonding
1. An ATOM has a NEUTRAL charge when it has the SAME number of PROTONS and
ELECTRONS
2. ION  is an ATOM that has TRANSFERRED (GAINED or LOST) one or more of its VALENCE
electrons giving the ION either a POSITIVE (“+“) charge or NEGATIVE (“-“)
ELECTRICAL charge making it NO longer NEUTRAL
a. Since an atom can NOT CHANGE its PROTON count, IONS form when one or more of the
atom’s VALENCE electrons are TRANSFERRED (GAINING or LOSING electrons)
*b. The number of PROTONS and NEUTRONS of an ION do NOT CHANGE
*1. Therefore, the ATOMIC number and ATOMIC MASS number of each ION stays the
SAME
*3. CATION  is a POSITIVELY charged ION
*a. Formed when atoms LOSE (by transferring) ELECTRON(S), giving the atom a GREATER
amount of POSITIVE charges in its NUCLEUS than NEGATIVE electrons
*b. (e.g.) sodium [Na]  Atomic number (PROTON count) = “11”; Electron count = “11”;
Group number = “1”; Valence Electron count = “1”;
Transfer (by LOSING) “1” electron = NEW electron count = “10”
sodium = CATION = Na+
(e.g.) boron [B]  Atomic number (PROTON count) = “5”; Electron count = “5”;
Group number = “13”; Valence Electron count = “3”;
Transfer (by LOSING) “3” electrons = NEW electron count = “2”
boron = CATION = B 3+
*4. ANION  is a NEGATIVELY charged ION
*a. Formed when atoms GAIN (by transferring) ELECTRON(S), giving the atom a GREATER
amount of NEGATIVE charges in its ORBITALS than POSITIVE charges in its NUCLEUS
*b. (e.g.) chlorine [Cl]  Atomic number (PROTON count) = “17”; Electron count = “17”;
Group number = “17”; Valence Electron count = “7”;
Transfer (by GAINING) “1” electron = NEW electron count = “18”
chlorine = ANION = Cl ((e.g.) oxygen [O]  Atomic number (PROTON count) = “8”; Electron count = “8”;
Group number = “16”; Valence Electron count = “6”;
Transfer (by GAINING) “2” electrons = NEW electron count = “10”
oxygen = ANION = O 2*5. GENERAL RULE for forming CATIONS and ANIONS is:
*a. Elements with “1 - 3” ELECTRONS in their VALENCE ORBITALS will TRANSFER by
LOSING electrons
*b. Elements with “5 - 7” ELECTRONS in their VALENCE ORBITALS will TRANSFER by
GAINING electrons
6. PROPERTIES of IONIC compounds
a. CHEMICAL bond  is the FORCE that holds two or more ATOMS or IONS together
b. IONIC BOND  is the FORCE that brings OPPOSITELY charged IONS together
1. IONIC BONDS form when ELECTRONS from a METAL atom (CATION) are
TRANSFERRED (LOSING and GAINING) to a NON-METAL atom (ANION)
2. ELECTRONS move FROM the VALENCE orbital of the METAL atom (CATION)
TO the VALENCE orbital of the NON-METAL atom (ANION) forming a “+” and “-”
ION pair
3. Bonded together the VALENCE orbitals of both ATOMS are now FILLED and a
NEUTRAL ionic COMPOUND is formed
*a. (e.g.) Mg2+ (CATION) + Cl2- (ANION)  2 MgCl (NEUTRAL)
c. IONIC COMPOUND  is a COMPOUND formed from an ION pair BONDING together
1. Ionic compound PROPERTIES include:
a. Large CRYSTAL LATTICE structures (REPEATING 3-D PATTERN)
1. CRYSTALS are hard and BRITTLE
b. Strong BONDS
c. High MELTING and BOILING POINT temperatures
*1. Usually CRYSTALLINE solid at ROOM temperature (250 C)
d. LIQUID ionic COMPOUNDS are GOOD conductors of ELECTRICITY
because the ATOMS are LOOSELY held together
1. SOLID ionic COMPOUNDS are poor CONDUCTORS of electricity
because the ATOMS are too TIGHTLY held together and the electrical
CHARGE can NOT pass through
2. Ionic COMPOUNDS DISSOLVED in water can CONDUCT electricity
e. Ionic COMPOUNDS are SOLUBLE (the ability to DISSOLVE) in water [H2O]
1. Water MOLECULES surround each ION and move them apart from
each other, which causes the separated IONS to DISSOLVE
*d. Naming an IONIC COMPOUND (established by the IUPAC):
*1. Ionic compounds bear the NAME of the POSITIVE ion (which is listed FIRST) followed
by the NAME of the NEGATIVE ion with the SUFFIX “-ide”
*a. (e.g.) lithium [Li]  Group “1”; VALENCE electron = “1”
CATION = transfer by LOSING “1” electron = “Li+”
chlorine [Cl]  Group “17”; VALENCE electron = “7”
ANION = transfer by GAINING “1” electron = “Cl - ”
*1. Li + (lithium) + Cl- (chlorine)  LiCl (LITHIUM CHLORIDE)
*b. (e.g.) magnesium [Mg]  Group “2”; VALENCE electrons = “2”
CATION = transfer by LOSING “2” electrons = “Mg2+”
bromine [Br]  Group “17”; VALENCE electrons = “7”
ANION = transfer by GAINING “1” electron = “Br - ” (for “2” Br atoms)
*1. Mg 2+ (magnesium) + Br2 - (bromine)  2 MgBr (MAGNESIUM BROMIDE)
*c. (e.g.) aluminum [Al]  Group “13”; VALENCE electrons = “3”
CATION = transfer by LOSING “3” electrons = “Al3+”
sulfur [S]  Group “16”; VALENCE electrons = “6”
ANION = transfer by GAINING “2” electron = “S 2- ”
*1. 2 Al 3+ (aluminum) + 3 S 2- (sulfur)  Al2S3 (ALUMINUM SULFIDE)
B. Covalent Bonds
1. COVALENT BONDS  is a FORCE that holds NON-METAL atoms together by SHARING one
or more PAIRS of ELECTRONS
a. The SHARED electrons fill the empty spaces in the VALENCE orbitals in each ATOM
*b. COVALENT bonds can form between two ATOMS of the SAME element or two
DIFFERENT atoms
*1. ELEMENTS that form COVALENT bonds:
*a. “HONC” RULE  hydrogen [H]; oxygen [O]; nitrogen[N]; carbon[C]
*b. Group 14  because all the ELEMENTS have “4” VALENCE electrons
*c. Group 17  only when the ELEMENTS bond with THEMSELVES making a
DIATOMIC MOLECULE
*2. DIATOMIC MOLECULE  is a MOLECULE that forms from the COVALENT
bonding between two ATOMS of the SAME element
*a. (e.g.) “2” hydrogen [H] atoms = H2; “2” oxygen [O] atoms = O2; “2” nitrogen
[N] atoms = N2; the HALOGENS (Group 17) form DIATOMIC molecules
*b. (e.g.) HALOGENS (Group 17): fluorine [F] = F2; chlorine [Cl] = Cl2;
bromine [Br] = Br2; iodine [I] = I2
2. MOLECULE  is a type of particle that is the SMALLEST unit of a COMPOUND held together
by chemical BONDS that have all the PROPERTIES of that COMPOUND, while
basically having NO “net” electrical CHARGE
*a. POLAR bond  a COVALENT bond in which the electrons are SHARED UNEQUALLY
*1. Some ATOMS pull more STRONGLY on the SHARED electrons than the others, pulling
them CLOSER to the NUCLEI of one atom causing the ATOMS to have slight
ELECTRICAL charges
*a. The ELECTRICAL charges are NOT as STRONG as the charges of IONS, but
enough to make one atom SLIGHTLY “-” and the other SLIGHTLY “+”
*b. (e.g.) WATER molecule = [H] atoms closer together (“+”); [O] atom more (“-”)
*b. NONPOLAR bond  a sharing of electrons pulling EQUALLY in OPPOSITE directions,
which CANCELS the POLAR bond
*1. (e.g.) CO2 molecule = [O] atoms more (“-”); [C] more (“+”); CANCEL each other
3. COVAENT COMPOUND  is a MOLECULE formed from the COVALENT BONDING of
NON-METAL atoms
a. Covalent COMPOUND PROPERTIES include:
1. AMORPHOUS solid structures
a. (e.g.) WAX, glass
2. Strong BONDS holding the ATOMS of the molecules together
a. WEAK force holding MOLECULES close to one another
3. Low MELTING and BOILING POINT temperatures
a. COVALENT bonds do NOT break, but rather the MOLECULES separate from
the other MOLECULES instead
*b. Usually a GAS or a volatile LIQUID (evaporates easily) at room temperature
(250 C)
4. Most COVALENT compounds are POOR conductors of ELECTRICITY because their
MOLECULES are NEUTRAL and are NOT charged like IONS
5. Most COVALENT compounds are INSOLUBLE (NOT able to be DISSOLVE) in
water [H2O]
*4. Naming a Molecular (Covalent) COMPOUND (established by the IUPAC):
*a. Covalent compounds use the NAME of the ELEMENT farthest to the LEFT on the periodic
table FIRST, (if from the SAME group use the ELEMENT with the HIGHER ATOMIC
number) followed by the NAME of the “2nd” element with the SUFFIX “-ide” and finally use
a PREFIX (mono- = 1; di- = 2; tri- = 3; tetra- = 4; penta- =5; hexa- =6; hepta- = 7; octa- = 8;
nona- = 9; deca- = 10 ) to represent the number of ATOMS of each ELEMENT
*1. (e.g.) carbon monoxide [CO] carbon [C] Group “14” and oxygen [O] Group “16”
carbon OXIDE = carbon MONoxide  CO
*a. Prefix “mono-” is ONLY used with the SECOND element
*2. (e.g.) carbon dioxide [CO2]  carbon [C] Group “14” and oxygen [O] Group “16”
carbon OXIDE = carbon DIoxide  CO2
*3. (e.g.) ammonia [NH3]  nitrogen [N] Group “15” and hydrogen [H] Group “1”
nitrogen HYDRIDE = nitrogen TRIhydride  NH3
(common name = AMMONIA)
*4. (e.g.) C2O  carbon [C] Group “14” and oxygen [O] Group “16”
carbon OXIDE = DIcarbon MONoxide  C2O
(also known as = COCONUT water)
*5. ELEMENTS that form COVALENT and IONIC bonds:
*a. from the “HONC” RULE  oxygen [O]; nitrogen[N] only
*b. Group 17  because they can form BONDS between two ATOMS of the SAME element
(covalent bond - DIATOMIC MOLECULES) or two DIFFERENT atoms (ionic bond)
*6. ELEMENTS that do NOT FORM any BONDS:
*a. GROUP 18 (the NOBLE or INERT gases)  because their VALENCE orbitals are FULL
C. Metallic Bonds
1. METALLIC BOND  is the BOND that forms between METAL atoms of the SAME element
when their VALENCE orbitals overlap
a. METALLIC bonds are WEAKER than IONIC or COVALENT bonds
2. METALLIC bonding is the SHARING of the FREE, “SEA” of electrons moving in ALL directions
around the “FIXED” positive IONS (which have TEMPORARILY lost their VALENCE electrons)
giving METALS their unique PROPERTIES
a. The POSITIVE metal IONS have enough valence ELECTRONS moving around them to
maintain the METALLIC BOND and hold the IONS together
3. METALLIC bonding PROPERTIES include:
a. High MELTING POINT temperatures
b. GOOD conductors of ELECTRICITY
1. VALENCE electrons of the METAL (e.g. copper [Cu]) are FREE to move because they
are NOT connected to any one ATOM
c. MALLEABILITY  ability of METALS to be RESHAPED by being hammered into
SHEETS or BENT without breaking
d. DUCTILE  ability to be RESHAPED into long, thin WIRES