UNIT 2 NOTES
59
Electronic Configuration
The location of an electron is described by 3 terms.
1st Term: Shell (n)
shell
2nd Term: subshell (sublevel)
- designated by s, p,d,f
( g,h,i )
- refers to the shape(s) of
the area in which the electron
can be located.
- also designates an energy level
within the shell.
The first shell (1) has one subshell (s).
The s subshell is spherical in shape and has 1 orbital
3rd Term: orbital
Each orbital can accommodate
2 electrons
subshell
- designated by s, p,d,f
- refers to the shape(s) of
the area in which the electron
can be located.
- also designates an energy level
within the shell.
- relative energy: s < p < d < f
s subshell: spherical
1 orbital
p subshell: pair of lobes
z
x
y
x
z
y
d subshell: double dumbells (note not required to know f or g
subshells or name)
x
y
xy
xz
yz
x2-y2
z2
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orbitals in subshells
Sublevel
s
p
d
f
# Orbitals # electrons
1
2
3
6
5
10
7
14
Remember shells are numbered 1,2,3, etc.
Three rules are used to build the
electron configuration:
– Aufbau principle
– Pauli Exclusion Principle
– Hund‘s Rule
Aufbau Principle
• Electrons occupy orbitals of lower
energy first.
Blocks in the Periodic Table
61
Aufbau
Diagram
Order of aufbau series shell filling
-Pauli Exclusion Principle
(Wolfgang Pauli, Austria, 1900-1958)
-Electron Spin Quantum Number
• An orbital can hold only two electrons and they
must have opposite spin.
Hund‘s Rule
In a set of orbitals, the electrons will fill the
orbitals in a way that would give the maximum
number of parallel spins (maximum number of
unpaired electrons).
Analogy: Students could fill each seat of a school
bus, one person at a time, before doubling up.
62
Aufbau
Diagram for
Boron
Aufbau
Diagram for
Nitrogen
Notations of Electron
Configurations
• Standard
• Shorthand
• Orbital notation/diagram
Standard Notation
of Fluorine
Number of electrons
in the sub shell 2,2,5
1s2 2s2 2p5
Subshells
63
Shorthand Notation
• Use the last noble gas that is located in the
periodic table right before the element.
• Write the symbol of the noble gas in
brackets.
• Write the remaining configuration after the
brackets.
• Ex: Fluorine: [He] 2s2 2p5
Electron Box Diagram Notation
• Draw out a box for each orbital in each
subshell
• Apply the rules for distributing electrons
Blocks in the Periodic Table
Our Orbital Picture of Be
The first shell (1s) is filled.
The 2s orbital has 2 e- present.
The 2p orbitals are empty.
The 2p orbitals have room for 6 e-
64
Our Orbital Picture of C
The first shell (1s) is filled.
The 2s orbital has 2 e- present.
2 2p orbitals have 1 e- each.
How to draw a Bohr Model. Examples below
65
Electron Configuration Practice Worksheet
In the space below, write the unabbreviated electron configurations of the following elements:
1)
sodium ________________________________________________
2)
iron
________________________________________________
3)
bromine
________________________________________________
4)
barium
________________________________________________
5)
neptunium
________________________________________________
In the space below, write the abbreviated electron configurations of the following elements:
6)
cobalt
________________________________________________
7)
silver
________________________________________________
8)
tellurium
________________________________________________
9)
radium
________________________________________________
10)
lawrencium
________________________________________________
Determine what elements are denoted by the following electron configurations:
11)
1s22s22p63s23p4 ____________________
12)
1s22s22p63s23p64s23d104p65s1 ____________________
13)
[Kr] 5s24d105p3 ____________________
14)
[Xe] 6s24f145d6 ____________________
15)
[Rn] 7s25f11 ____________________
Determine which of the following electron configurations are not valid:
16)
1s22s22p63s23p64s24d104p5 ____________________
17)
1s22s22p63s33d5 ____________________
18)
[Ra] 7s25f8 ____________________
19)
[Kr] 5s24d105p5 ____________________
20)
[Xe] ____________________
66
67
FLAME TEST LAB
In this activity, you will investigate the colors of flame produced by solutions of metal salts.
When a substance is heated in a flame, the atoms absorb energy from the flame. This absorbed energy allows the
electrons to be promoted to excited energy levels. From these excited energy levels, there is a natural tendency for
the electrons to make a transition or drop back down to the ground state. When an electron makes a transition from
a higher energy level to a lower energy level, a particle of light called a photon is emitted. Both the absorption and
emission of energy are quantized – only exact amounts of energy are required.
An electron may drop all the way back down to the ground state in a single step, emitting a photon in the process.
Alternatively, an electron may drop back down to the ground state in a series of smaller steps, emitting a photon
with each step. In either case, the energy of each emitted photon is equal to the difference in energy between the
excited state and the state to which the electron relaxes. The energy of the emitted photon determines the color of
light observed in the flame.
In this activity, q-tips dipped in solutions of metal salts are heated using a Bunsen burner, producing different
colored flames. By comparing the color given off by an unknown with the known metal salts, the identity of the metal
salt can be determined.
Flame Tests Activity
Materials:



Bunsen burner
Q-tips
Solutions (1.0 mol/L) of the following metal salts
 lithium chloride
 strontium chloride
 calcium chloride
 copper(II)chloride
 sodium chloride
 potassium chloride
 UNKNOWN
Procedure:
1. Light the Bunsen burner and open the air vent to obtain a non-luminous flame with two blue cones. Be sure to
avoid a yellow flame. (Why?)
2. A each station. Carefully place the end of q-tip that was soaked in the metal salt solution at the top of the inner
blue cone. Record the color and intensity (bright/faint) of the flame in the data table. The color given off by the
salt is the initial color observed, not the yellow-orange color produced by the burning q-tip. (To avoid burning
the q-tip, wave it through the flame rather than holding it right in the flame).
3. Repeat with the other 7 stations. Be sure to record the colors as ―precisely‖ as possible.
Data Table:
Metal found in the
salt
Flame Color and Intensity
Lithium
Strontium
68
Calcium
Copper
Sodium
Potassium
UNKNOWN
Identity:_____________
69
Flame Test Analysis
1. List the colors observed in this lab from the highest energy to the lowest energy.
2. List the colors observed in this lab from the highest frequency to the lowest frequency.
3. List the colors observed in this lab from the shortest wavelength to the longest wavelength.
4. What is the relationship between energy, frequency, and wavelength?
5. Do you think we can use the flame test to determine the identity of unknowns in a mixture? Why
or why not?
6. How are electrons ―excited‖ in this part of the experiment? What does it mean when the electrons
are ―excited‖?
7. Explain why we did not see distinct lines (like on an emission spectrum) when the metal salts
were burned.
8. Why do different chemicals emit different colors of light?
9. Why do you think the chemicals have to be heated in the flame first before the colored light is
emitted?
10. Colorful light emissions are applicable to everyday life. Where else have you observed colorful
light emissions? Are these light emission applications related? Explain.
70
Development of a New Atomic
Model
Connections…
• Last time- Democritus, Thomson, Dalton,
Rutherford
• In 1900s, new model evolved as a result of
studying light absorbed and emitted by matter
Wave Description of Light
• Electromagnetic radiation (ER): form of
energy that exhibits wavelike behavior as it
travels through space
• All forms of ER together make the
electromagnetic spectrum
71
• All forms of ER move at a constant speed of
about 3.0 x108 m/s (speed of light, c)
• Wavelength (λ): distance between
corresponding points on adjacent waves.
• Frequency (ν): number of waves that pass a
specific point in a given time, usually one
second.
• Unit: Hertz (Hz), aka (1/s) or (s-1)
• For electromagnetic radiation, frequency and
wavelength are related
C=λν
• If λ increases, what must happen to ν? Does c
change?
Photoelectric Effect
• Photoelectric effect: emission of electrons
from a metal when light shines on that metal
• Experiment: no e- were emitted if light’s
frequency was below a certain threshold,
regardless of time shone.
• According to wave theory, shining any light
long enough should supply enough energy to
eject an e-
Particle description of light
• Max Planck, 1900s, suggested that object emit
energy in small, specific amounts called
quanta
• Quantum: minimum amount of energy that
can be gained or lost by an atom
72
• Planck proposed a relationship between a
quantum of energy and the frequency of
radiation
E=hν
• E is energy, in Joules, of a quantum of
radiation
• h is Planck’s constant (fundamental physical
constant)= 6.626 x 10-34 J*s
• ν is freqency of radiation
• 1905 Einstein expands on this idea. ER have
dual wave/particle nature.
• While light emits many wavelike particles, it
can also be thought of as a stream of those
particles
• Einstein named the particles photons
• Photon: particle of ER having zero mass and
carrying a quantum of energy
• Energy of a particular photon depends on the
frequency of the radiation
Ephoton= hν
• Einstein’s Explanation: ER is absorbed by matter
only in whole numbers of photons
– For e- to be ejected, must be struck by single photon
possessing at least minumum energy
– According to equation, this energy corresponds to
frequency
– If photon’s frequency is below minimum, no e- ejected
The Hydrogen Atom
• Ground state: lowest energy state of an
atom
• Excited state: atom has higher potential
energy than ground state
• When an excited atom returns to ground
state, it gives off energy it gained in the
form of ER.
• Production of colored signs (neon) is
example
73
Experiment
• Pass electrical current through H gas in
vacuum tube at low pressure
• Emits characteristic pink glow
• When the light was passed through prism,
separated into series of specific frequencies
and therefore wavelength (… equation?) of
visible light
• These bands are Hydrogen’s line emission
spectrum
Problem?
• Classical theory predicted that H atoms would
be excited by whatever amount of energy was
added to them.
• Expected to observe continuous range of
frequencies of ER, or continuous spectrum.
• Why had H only given off specific frequencies
of light?
•  Quantum Theory
74
• Excited H atom falls back from excited state to
ground state, and emits photon
• This energy is equal to difference between
initial and final state
Implications
• Since H atoms emit only specific frequencies,
difference between energy states must be
fixed.
• Therefore, e- of H atom exists only in very
specific energy states
• In 1913, Bohr proposed a model that linked
the atom’s electron with photon emission
• Energy is higher in orbits farther from
nucleus (like a ladder)
• Based on the wavelengths of hydrogen’s line
emission spectrum, Bohr calculated energies
the e- would have in the allowed energy level
for H atom.
• Bohr’s calculated values agreed with
experimentally observed values for lines in
each series
• Scientists tried to apply this model to other
element’s atoms
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Energy Calculations:
1.
What is the energy of a mole of photons with a wavelength of 1.60 10-3 m?
2.
What is the energy of a photon with a wavelength of 2.65 10-4 m?
3.
What is the frequency of a photon with a wavelength of 3.70 10-6 m?
4.
What is the frequency of a photon with a wavelength of 8.60 103 nm?
5.
What is the energy of a photon with a wavelength of 6.55 102 nm?
6.
What is the wavelength of a photon with an energy of 3.40 10-18 J?
7.
What is the wavelength of a photon with a frequency of 2.35 1012 s-1?
8.
What is the energy of a photon emitted by the 3 to 1 transition in a hydrogen atom?
a.
What is the wavelength of the photon emitted, in meters and nanometers?
76
GENERAL PROPERTIES OF THE PERIODIC TABLE
Mendeleev's Periodic Table (1869)
A. Organization
1. Vertical columns in atomic weight order
a. Mendeleev placed elements in rows with similar properties
2. Horizontal rows have similar chemical properties
B. Missing Elements
1. Gaps existed in Mendeleev‘s table
a. Mendeleev predicted the properties of the ―yet to be discovered‖ elements
(1) Scandium, germanium and gallium agreed with his predictions
C. Unanswered Questions
1. Why didn't some elements fit in order of increasing atomic mass?
2. Why did elements exhibit periodic behavior?
Moseley and the Modern Periodic Table (1911)
A. Protons and Atomic Number
1. The periodic table was found to be in atomic number order, not atomic mass order
B. The Periodic Law
1. The physical and chemical properties of the elements are periodic functions of their atomic numbers
2. Elements with similar properties are found at regular intervals within the periodic table
* Moseley was killed in battle in 1915, during WWI. He was only 28 years old
Organization of the Table
1. Groups or Families
a. Vertical columns containing elements with similar chemical properties
2. Periods (series)
a. Horizontal rows of elements
3. Metals and Nonmetals
a. A stair-step line on the table separates the metals from the nonmetals
b. Metalloids (Semimetals) straddle the line and have properties of both metals and nonmetals
4. Lanthanide and Actinide Series (Inner Transition Metals)
a. Metals and man-made metal elements
5. Group 1 – Alkali metals (the most reactive metal elements) (except hydrogen (H) also in this group)
6. Group 2 – Alkaline earth metals (very reactive metal elements)
7. Group 17 – Halogens (the most reactive nonmetal elements)
8. Group 18 – Noble gases (the least reactive elements – inert and very stable)
Types of Elements
A. Metals
1. Luster
2. Good conductors of heat and electricity
3. Malleable
4. Ductile
5. High tensile strength
B. Nonmetals
1. Many nonmetals are gases at room temperature
2. Solid nonmetals tend to be brittle and non-lustrous
3. Poor conductors of heat and electricity
C. Metalloids
1. Some properties of metals and some properties of nonmetals
2. Solids at room temperature
3. Semiconductors of electricity
D. Noble Gases
1. All are gaseous members of group 18
2. Generally unreactive and stable
77
Trends in the
Periodic Table
Periodic Table
• Prior to 1860 no agreement/method to
accurately determine masses of atoms.
• First International Congress of Chemists –
1860
– Stanislao Cannizzaro presented method for
accurately measuring atomic masses
– Looked for relationships between atomic
masses and other properties of elements
• First tables arranged elements by atomic
weight
– Could not agree on atomic weights therefore
tables were different
John Newlands
• Noticed elements properties repeated
every 8th element when arranged by
atomic mass
• Named this phenomenon ―the Law of
Octaves‖
• Did not work for all elements
78
Julius Lothar Meyer
• Developed first modern
table
– Consisted of 28 elements
divided into 6 families
– Families (groups) had similar
chemical and physical
properties
– Discovered all elements in
same family had same
number of valence e- -outermost electrons in
highest energy level
– Why?
Dmitri Mendeleev
• Noticed that properties
repeat themselves at
certain intervals
• Arranged all known
elements into one table
based on properties– 1869
• 1871 - Proposed the
―Periodic Law‖
• Based on the properties
spaces were left for
unknown elements (Sc,
Ga, Ge)
79
• Upon discovery of other elements
inconsistencies were found with
Mendeleev’s table
• Atomic masses improved and they no
longer arranged the elements by
increasing atomic mass
• Why can most elements be arranged by
atomic mass?
• What was the reason for chemical
periodicity?
Henry Mosely
• Discovered elements contain unique
number of protons (atomic number) - 1911
• Arranged elements by atomic number 1913
• Fully explained the Periodic Law
Periodic Law
• The physical and chemical properties of
the elements are periodic functions of their
atomic numbers.
• Aka – when elements are arranged by
increasing atomic number, elements with
similar properties appear at regular
intervals.
80
Parts….
•
•
•
•
•
Alkali metals – group 1
Alkaline earth metals – group 2
Halogens – group 17
Transition metals – d block elements
Inner Transition metals
– Lanthanides (elements 58-71) added in early
1900’s
• Have very similar properties
– Actinides (elements 90-103)
s-Block Elements
• Groups 1 & 2
• All elements in group 1 & 2 will have an
electron configuration of
– ns1 or ns2 where n = highest energy level
occupied
Alkali Metals
• Group 1 elements
• In the elemental state
– Soft
– Silvery metal
– High melting points
– Extremely reactive therefore are not found in
elemental state in nature
• React violently with water to produce
hydrogen gas
Alkaline – Earth Metals
• Group 2 elements
• Outer most s orbital is full
– Do not exhibit stability (outer p orbital is
empty)
• Properties
– Harder, denser than group 1
– Higher melting points than group 1
– Not as reactive but too reactive to be found in
nature in elemental form
81
Burning Mg
Hydrogen & Helium
• H has same valence electrons as group 1
but does not share any other properties
• He share same electron configuration
(valence e-) as group 2 but does not share
same properties
– Placed with group 18 because it is very stable
d-block elements
• Transition elements
– Beginning filling the 3d orbitals
– Good conductors of electricity
– High luster
– Less reactive than s-block elements
• Can be found in elemental form
Exceptions in the d-block
• The following elements have odd
configurations
– Cr: [Ar]4s13d5
– Cu: [Ar]4s13d10
– Ag: [Kr]5s14d10
• More stable with half filled s & d orbitals or
full d orbital
• Exceptions follow throughout the d
element similar to Chromium and Copper
82
p-block elements
• All elements in p block have a full s orbital
• Properties
– Contain all non metals except H & He
– Contain all metalloids (exhibit properties of both
metals and non metals)
• Have semi conducting properties
– Contains 6 metals
• Elements in s & p block make up the
representative elements
Halogens
• Group 7A/17
– Most reactive non metals (Fluorine is most
reactive)
– Will bond with a metal to form a salt
– F & Cl are gases at room temp
– Br is a liquid at room temp
– I & At are solids at room temp
Atomic Radius
• Atomic radius is simply the radius of the atom,
an indication of the atom's volume.
• Atomic radius is one-half the distance between
the two nuclei in a molecule consisting of two
identical atoms.
Trends in Atomic Size
• - Period - atomic radius decreases as you go
from left to right across a period.
• Why? Stronger attractive forces in atoms (as you
go from left to right) between the opposite
charges in the nucleus and electron cloud cause
the atom to be 'sucked' together a little tighter.
83
Trends in Atomic Size cont.
• Group - atomic radius increases as you go down a
group.
Why?
• There is a significant jump in the size of the nucleus
(protons + neutrons) each time you move from period to
period down a group.
• Additionally, new energy levels of elections clouds are
added to the atom as you move from period to period
down a group, making the each atom significantly more
massive, both is mass and volume.
Electronegativity
• Electronegativity is an
atom's 'desire' to grab
another atom's electrons.
Electronegativity Trends
• Period - electronegativity increases as you go from left
to right across a period.
• Why? Elements on the left of the period table have 1 -2
valence electrons and would rather give those few
valence electrons away (to achieve the octet in a lower
energy level) than grab another atom's electrons. As a
result, they have low electronegativity. Elements on the
right side of the period table only need a few electrons to
complete the octet, so they have strong desire to grab
another atom's electrons.
84
Electronegativity Trends cont.
• Group - electronegativity decreases as you go down a
group.
• Why? Elements near the top of the period table have few
electrons to begin with; every electron is a big deal. They
have a stronger desire to acquire more electrons.
Elements near the bottom of the chart have so many
electrons that loosing or acquiring an electron is not as
big a deal. This is due to the shielding effect where
electrons in lower energy levels shield the positive
charge of the nucleus from outer electrons resulting in
those outer electrons not being as tightly bound to the
atom.
Ionization Energy
• Ionization energy is the amount of
energy required to remove the
outermost electron/s.
• Ionization energy is closely
related to electronegativity.
Ionization Energy Trends
• Period - ionization energy increases as you go from left
to right across a period.
• Why? Elements on the right of the chart want to take
others atom's electron (not given them up) because they
are close to achieving the octet. The means it will require
more energy to remove the outer most electron.
Elements on the left of the chart would prefer to give up
their electrons so it is easy to remove them, requiring
less energy (low ionization energy).
85
Ionization Energy Trends cont.
• Group - ionization energy decreases as
you go down a group.
• Why? The shielding effect makes it easier
to remove the outer most electrons from
those atoms that have many electrons
(those near the bottom of the chart).
Reactivity
Reactivity refers to how likely or
vigorously an atom is to react with
other substances.
This is usually determined by two
things:
1) How easily electrons can be
removed (ionization energy)
from an atom
86
2) or how badly an atom wants
to take other atom's electrons
(electronegativity)
The transfer/interaction of
electrons is the basis of chemical
reactions.
Reactivity of Metals
Period - reactivity decreases as you go from
left to right across a period.
Group - reactivity increases as you go down a
group
Why? The farther to the left and down the
periodic chart you go, the easier it is for
electrons to be given or taken away, resulting in
higher reactivity.
Reactivity of Non-Metals
Period - reactivity increases as you go from
the left to the right across a period.
Group - reactivity decreases as you go down
the group.
Why? The farther right and up you go on the
periodic table, the higher the
electronegativity, resulting in a more
vigorous exchange of electron.
87
Ionic Radius vs. Atomic
Radius
• Metals - the atomic radius of a metal is generally
larger than the ionic radius of the same element.
• Why? Generally, metals loose electrons to achieve the
octet. This creates a larger positive charge in the
nucleus than the negative charge in the electron cloud,
causing the electron cloud to be drawn a little closer to
the nucleus as an ion.
Ionic Radius vs. Atomic Radius
cont.
Non-metals - the atomic radius of a nonmetal is generally smaller than the ionic
radius of the same element.

Why? Generally, non-metals loose
electrons to achieve the octet. This creates
a larger negative charge in the electron
cloud than positive charge in the nucleus,
causing the electron cloud to 'puff out' a little
bit as an ion.

Ionic Radius vs. Atomic Radius
Summary of Periodic Trends
88
Periodic Table Trend War
Purpose: To teach students periodic trends regarding the properties of elements.
Materials: 1 die, index cards with main group element symbols written on them (one
element per card), a list of trends written on the board numbered from 1 - 6. Trends
should include such things as ionization energy, electronegativity, atomic radius, ionic
radius, etc… A "wild" category can be included which allows the roller of the dice to
choose the trend.
Playing the game:
• Students are placed in groups of 4 - 5.
• Cards are dealt until each student has the same number of cards. Each student
should have about 7 cards per hand.
• The dealer begins play by throwing the die. The number of the die determines
the trend being played.
• After the trend for the hand is determined the dealer plays the first card. Play
continues to the left of the dealer. The card with the highest value for the
current trend wins. The player who takes the hand rolls the die and makes
the next lead. Multiple hands may be played.
89
Common Ions
H+
Li+
Na+
K+
NH4+
Ag+
1+ ions
Hydrogen
Lithium
Sodium
Potassium
Ammonium
Silver
Mg
Ca2+
Ba2+
Zn2+
2+ ions
Magnesium
Calcium
Barium
Zinc
Al3+
3+ ions
Aluminum
2+
-
H
FClBrINO2NO3BrO3ClOClO2ClO3ClO4IO3OHCNHCO3HSO3HSO4C2H3O2MnO4-
1- ions
Hydride
Fluoride
Chloride
Bromide
Iodide
Nitrite
Nitrate
Bromated
Hypochlorite
Chlorite
Chlorate
Perchlorate
Iodate
Hydroxide
Cyanide
Bicarbonate
Bisulfite
Bisulfate
Acetate
Permanganate
O2S2CO32SO32SO42CrO42S2O32C2O42O22-
2 - ions
Oxide
Sulfide
Carbonate
Sulfite
Sulfate
Chromate
Thiosulfate
Oxalate
Peroxide
3-
N
P3PO33PO43BO33-
3- ions
Nitride
Phosphide
Phosphite
Phosphate
Borate
Variable Charge Metals
Fe2+
Fe3+
Latin Names
Ferrous
Ferric
Cu+
Cu2+
Cuprous
Cupric
Pb2+
Pb4+
Plumbous
Plumbic
Sn2+
Sn4+
Stannous
Stannic
Hg2 2+ (1+)
Hg2+
Mercurous
Mercuric
90
Compounds and Molecules
A molecule is made up of two or more atoms chemically combined.
A molecule is the smallest part of an element or compound that can exist on its own
Compounds are substances made from two or more different elements chemically combined.
So what does this phrase ‗chemically combined‘ mean?
First we need to note that there are two ways for atoms to combine chemically, and that another word for
the process of combining chemically is ‗bonding‘.
Bonding occurs because all atoms try to have Bonding
a full outer shell, and will lose, gain or share electrons
in order to do so
There are two separate methods by which atoms join up:
1. Ionic bonding occurs when an atom loses or gains electrons
2. Covalent bonding occurs when an atom shares electrons.
Ionic bonding
An ion is a charged atom – i.e. an atom which has lost or gained electrons
If an atom loses an electron then it becomes a positive ion.
If an atom gains an electron then it becomes a negative ion.
In ionic bonding positive ions are attracted to negative ions.
An ionic bond is a bond formed by the force of attraction between two oppositely charged ions
An example of an ionic bond is sodium chloride
The bonds in sodium chloride are formed by sodium atoms losing electrons and chlorine atoms gaining
electrons.
Examples of ionic bonding: (i) Sodium Chloride (NaCl)
The atomic number of sodium is 11.
It has an electronic configuration of 2, 8, 1. This means that it needs to lose the one electron which it has
in its outer shell in order to have a complete outer shell.
The atomic number of chlorine is 17.
Chlorine has an electronic configuration of 2, 8, 7. This means
it needs to gain one electron in order to have a complete outer
shell.
When a sodium atom bonds with a chlorine atom the sodium
atom loses its outer electron to form a positive ion while the
chlorine atom gains an electron to form a negative ion.
91
The two atoms now have opposite charges and because opposite charges attract both atoms move off
together as a sodium-chloride (NaCl) molecule.
Examples of ionic bonding: (ii) Magnesium Oxide
(MgO)
The atomic number of magnesium is 12.
It has an electronic configuration of 2, 8, 2. This means that
it needs to lose the two electrons which it has in its outer
shell in order to have a complete outer shell.
The atomic number of oxygen is 8.
Oxygen has an electronic configuration of 2, 6. This means
it needs to gain two electrons in order to have a complete
outer shell.
When a magnesium atom bonds with an oxygen atom the
sodium atom loses its two outer electrons to form a positive
ion while the oxygen atom gains two electrons to form a
negative ion.
The two atoms now have opposite charges and because
opposite charges attract both atoms move off together as a
magnesium-oxide (MgO) molecule.
Covalent bonding
A covalent bond is a bond formed by different atoms sharing electrons in order to have a
complete outer shell
Examples of covalent bonds are H2, O2, H2O and CH4 (methane) molecules.
The bonds in these molecules are formed by the atoms in the molecules sharing electrons with each
other..
Examples of covalent bonding: (i) A hydrogen molecule (H2)
The atomic number of hydrogen is 1.
It needs to have two electrons in its outer shell, and so if it ‗bonds‘ with
another hydrogen atom and they both ‗share‘ their electron with each other, it
has the effect of allowing both atoms to have a complete outer shell.
Examples of covalent bonding: (ii) An oxygen molecule (O2)
The atomic number of oxygen is 8.
It has an electronic configuration of 2, 6.
It needs to have 8 electrons in its outer shell, and so if it ‗bonds‘ with another
oxygen atom and they both ‗share‘ two of their electrons with each other, it
has the effect of allowing both atoms to have a complete outer shell.
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Examples of covalent bonding: (iii) A water molecule (H2O)
The atomic number of hydrogen is 1 so it has 1 electron in its outer shell.
The atomic number of oxygen is 8 so it has 6 electrons in its outer shell.
Oxygen can share 1 electron with one hydrogen atom and a second electron with a second hydrogen atom,
which has the effect of allowing all atoms to have a complete outer shell.
Examples of covalent bonding: (iv) A methane molecule (CH4)
The atomic number of hydrogen is 1 so it has 1 electron in its outer shell.
The atomic number of carbon is 6 so it has 4 electrons in its outer shell.
Carbon can share 1 electron with each of 4 different hydrogen atoms,
which has the effect of allowing all atoms to have a complete outer shell.
Properties of Ionic and Covalent compounds
Ionic Compounds
Usually solid at room temperature
High melting and boiling points
Soluble in water
Conducts electricity
Undergo fast reactions
Covalent Compounds
Usually liquid or gas at room temperature
Low melting points
Insoluble in water
Do not conduct electricity
Undergo slow reactions
WHAT ARE METALS?
 A metal is an element that readily forms positive ions (cations) and has metallic bonds.
 On the periodic table, a diagonal or stair step line drawn from boron (B) to polonium (Po)
separates the metals from the nonmetals. Elements on this line are metalloids, sometimes called
semi-metals; elements to the lower left are metals; elements to the upper right are nonmetals.
 Almost 80% of the elements on the periodic table are metals.






METALLIC BONDS
The valence electrons of metal atoms can drift freely from
one part of the metal to another- this is sometimes called
a ―sea of electrons‖
Metallic bonds consist of the attraction between these
free floating electrons and the positively charged metal
ions (cations). This attraction is the “bond” that holds
metals together.
PHYSICAL PROPERTIES OF METALS
Lusterous- they are shiny!
High density- atoms are tightly packed.
Good conductors of electricity
o Reason- electrons can flow freely.
Ductile- they can be drawn into wires AND Malleable- they can be hammered into shapes
o Reason- cations can slide easily past each other because the sea of electrons
insulates them and prevents strong repulsions.
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CHEMICAL PROPERTIES OF METALS
 Most metals are chemically unstable and will react will oxygen in the air- that is they form oxides- over
varying timescales (for example iron rusts over years and potassium burns in seconds).
o The alkali metals react quickest followed by the alkaline earth metals.
o The transition metals take much longer to oxidize (such as iron, copper, zinc, nickel). Others, like
palladium, platinum and gold, do not react with the atmosphere at all.
o Some metals form a barrier layer of oxide on their surface which cannot be penetrated by further
oxygen molecules and thus retain their shiny appearance and good conductivity for many decades
(like aluminium, some steels, and titanium).


ALLOYS
Very few metals that you encounter daily are pure metals. Most metals are alloys, a mixture
of two or more elements of which at least one is a metal.
o Examples:
 Brass is an alloy of copper and zinc
 Sterling silver is an alloy of silver (92%) and copper (8%)
 Stainless steel is an alloy of iron (81%), chromium (18%), nickel (1%), and
trace amounts of carbon.
Alloys are important because their properties are often superior to those of their component
elements.
o Examples:
 Sterling silver is harder and more durable than pure silver, but still soft
enough to make jewelry and tableware.
 Brass is harder and easier to shape than either copper or zinc
Practice Questions
1. Name one property of ionic and one property of covalent compounds:
2. Draw a diagram (using suitable circuit symbols) of the circuit needed to investigate the ability of
ionic and covalent compounds to conduct electricity.
3. How could you tell whether a compound was ionic or covalent?
4. Draw a diagram showing the ionic bonding in the compound magnesium oxide.
5. Draw a diagram showing the bonding in a water molecule.
6. Draw a labelled diagram to show the type of bonding in an oxygen molecule.
7. Name the type of bonding present in a crystal of sodium chloride.
Give one property of this type of bonding:
8. In a negatively charged ion the number of
.
is greater than the number of
9. In a positively charged ion the number of
.
is greater than the number of
10. Underline the two elements in the list below that normally exist as molecules:
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Hydrogen
I.
II.
Chlorine
Helium
Aluminium
Gold.
Draw Lewis dot structures for the following elements and ions
A.
Mg
B.
Ga
C.
S
D.
Sn
E.
P
F.
I
G.
Rb
H.
Xe
I.
Cr
J.
Y
K.
S1-
L.
Ga 2+
Draw Lewis structures for the following molecules [single bonds]
A.
CH4
B.
BH3
C.
SCl2
D.
ICl
E.
PF3
F.
CH3OH
G.
NH2OH
H.
PBr3
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Greek Prefixes
Nomenclature (inorganic)Greek prefixes:
-Nonmetal
+ Nonmetal
1. First element retains its name. 3-8
2. Second element gets -ide ending. 4-9
3. Use Greek prefixes to identify the # of atoms 5-10
1–
2345-
678910½-
Examples N2O = ______________________
SO3 = ______________________
Now Try These:
carbon dioxide ___________________ BrF3 ________________________
diphosphorus pentasulfide __________ CS2 ________________________
1. Metal retains its name. • make sure ions join to form a neutral compound •
2. Non-metal retains its ionic name.
+
Examples: Na + Cl ---> ___________ name = __________________________
+2
Ca
-
+ Cl ---> ___________ name = __________________________
+
---> ___________ name = ___________________________
+
---> ___________ name = ___________________________
+
---> ___________ name = ___________________________
Now Try These:
lithium bromide ____________ Al2S3___________________________
magnesium hydroxide____________ (NH4)2SO3 _______________________
aluminum acetate ____________ Ba(NO3)2 _________________________
-variable charge metal + nonmetal
1a. Metal gets its charge written w/ Roman # in ( ) after name (w/ new system), -or1b. Metal gets Latin name (w/ old system).
2. Non-metal retains its ionic name.
+2
Examples: Fe + Cl ---> ________ name = __________________ (
)
+3
Fe
-
+ Cl ---> ________ name = __________________ (
+
---> ________ name = ____________________
+
---> ________ name = ____________________
)
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Now Try These:
manganese (IV) chloride __________ Cr(OH)3 ______________________
nickel (III) carbonate __________ CuCO3 ______________________
- Acids (positive ion = ―H+‖
1. Drop the ending on the negative ion.
The -ate ending changes to -ic acid.
The -ite ending changes to -ous acid.
The -ide ending uses the prefix hydro- and the suffix -ic acid
Examples: H2CrO4 = ______________________
HNO2 = ______________________
HCl = ______________________
Now Try These:
hydrobromic acid __________ H3PO3 _____________________
perchloric acid ____________ HI __________________________
Example: MgSO4 • 3 H2O = ___________________
Now Try These: calcium phosphate tetrahydrate = _________________
NaC2H3O2 • 2 H20 = ________________
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Nomenclature
1. carbon dioxide _________
2. potassium cyanide _________
3. selenium disulfide __________________
4. potassium chlorate _________
5. nitrous acid _________
6. zinc sulfate _________
7. aluminum acetate _________
8. copper(II) phosphate _________
9. disilicon trioxide _________
10. chloric acid _________
11. sodium chloride _________
12. aluminum iodide _________
13. barium cyanide _________
14. carbon disulfide _________
15. strontium nitrate _________
16. cuprous phosphate Cu3PO4
17. phosphorous acid _________
18. potassium hydroxide _________
19. bromine heptafluoride _________
20. lead (II) sulfide _________
21. carbon monoxide _________
22. ammonium acetate _________
23. mercuric borate _________
24. calcium hydride _________
25. boron trichloride _________
26. oxalic acid _________
27. BrF3 __________________
28. Li2C2O4 __________________
29. Fe3(PO4)2
30. SCl4 __________________
31. KHCO3 __________________
32. SnI2 __________________
33. HF hydrofluoric acid
34. PO3 __________________
35. PO3-3 __________________
36. CaCO3 __________________
37. Fe(IO3)2 __________________
38. CuCO3 copper (II) carbonate
39. CaF2 __________________
40. HNO3 __________________
41. (NH4)2S __________________
42. SO3 __________________
43. KNO3 __________________
44. Sn3(PO4)2 __________________
45. MgS2O3 __________________
46. Ca2C __________________
47. H2S __________________
48. CCl4 __________________
49. NaHSO3 __________________
50. NH4OH __________________
51. H3BO3 __________________
52. V(BrO3)5 __________________
53. MgSO4 • 7 H2O ______________________________
54. sodium acetate pentahydrate __________________
55. CuCl26 H2O __________________________________________
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Chemistry Unit #2 Review:
1.
Define a family. _______________________________________________________
2.
What is a period? ________________________________________________________
3.
What is the symbol for the following elements.
a. Magnesium _____________
b. Potassium ______________
What are the names of the following elements.
a. C __________________
b. Cl _________________
What period are the following elements in?
a. He _______________
b. Ge _________________
What group are the following elements?
a. Sulfur _______________
b. Ca _________________\
Give me an atom with the following characteristics.
a. Halogen _________________ b. Nonmetal ________________
c. Alkali metal ______________ d. metalloid ________________
e. Lanthanide series __________ f. Alkaline Earth metal ________________
g. Transition metal ___________ h. Nobel gas ________________
Write the electron longhand and shorthand configuration for
a. Li _______________________________________________________________
4.
5.
6.
7.
8.
9.
b.
Na ______________________________________________________________
c.
K _______________________________________________________________
What are valence electrons? __________________________________________________
_________________________________________________________________________
10. How many valence electrons are in the following element?
a. F ________
e. O ________
11.
12.
13.
14.
15.
16.
b. Cl ___________
c. Br ____________
d. I _____________
f. S ___________ g. Se ____________
h. Te ____________
On the blank periodic table below
Label the s, p, d, and f block elements
Create a circle that fills the whole box where the largest atom exists in the periodic table.
Put a dot where the smallest atom is in the periodic table.
Put a triangle on the box with the atom with the highest electronegativity
Put a square in the box with the lowest ionization energy
Label with arrows the trends for: atomic radius, ionic radius (metals and nonmetals), ionization energy, and
electronegativity
Name the following ionic compounds:
17.
NaBr __________________________________________
18.
CaO __________________________________________
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19.
Li2S __________________________________________
20.
MgBr2 __________________________________________
21.
Be(OH)2 __________________________________________
Write the formulas for the following ionic compounds:
22.
potassium iodide __________________________________________
23.
magnesium oxide __________________________________________
24.
aluminum chloride __________________________________________
25.
sodium nitrate __________________________________________
Write the names of the following covalent compounds:
26.
BF3 __________________________________________
27.
P2Br4 __________________________________________
28.
SF6 __________________________________________
29.
NH3 __________________________________________
30.
NO2 __________________________________________
Write the formulas of the following covalent compounds:
31.
nitrogen trichloride __________________________________________
32.
boron carbide __________________________________________
33.
dinitrogen trioxide __________________________________________
34.
phosphorus pentafluoride ______________________________________
35. Make a chart that compares and contrasts the different characteristics of ionic and covalent bonds.
36 What is the different between a molecular and empirical formula? Give an example.
From Electron Configuration Notes:
37. What is the shape of the s orbital? ________________ p orbital?______________ d orbital?
_________________
38. Which of the following orbitals is closest to the nucleus?
a. 2s
b. 3p
c. 1s
d. 4d
39. In the wave-mechanical (quantum) model of the atom, orbitals are regions of the most
probable locations of:
a. protons
b. positrons
c. neutrons
d. electrons
40. (CHALLENGE) Heiseinberg’s Uncertainty Principle states:
____________________________________________________________________________
____________________________________________________________________________/
100
41. Identify the following atom, 1s22s22p63s23p64s23d6 _________________________
42. Write out the orbital notation for F, N, Zn:
43. Write the electron configuration for the above elements:
44.Write the shorthand (noble gas) notation for the above elements:
F:_________________________________________________
F:__________________________________________
Zn:________________________________________________
Zn:_________________________________________
Sr:________________________________________________
Sr:_________________________________________
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