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Email: [email protected]: f-228Phone:
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CHAMPLAIN COLLEGESAINT-LAMBERT
GENERAL CHEMISTRY
202-NYA-205
Winter 2012
Joel Robichaud
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Course Outline
nomenclature empirical & molecular formulas stoichiometry
gas laws molarity
UNIT 1:Basics
UNIT 2:Atomic Theory
history of atomic theory the Bohr atom the modern approach
(quantum theory) quantum numbers electron configurations electron
affinity
UNIT 3:Periodicity &
Chemical Reactions
electron configuration & chemical properties of elements
ionization energy atomic and ionic size electronegativity &
electron affinity reactions of the main group elements writing
molecular & net ionic equations
UNIT 4:Chemical Bonding
analysis of ionic & covalent bonding writing Lewis
structures, resonance structures formal charges shapes of molecules
bond angles bond polarity dipole moments hybridization theory
orbital diagrams
UNIT 5:Intermolecular Forces
intermolecular forces intramolecularbonds dispersion forces,
dipole-dipole forces hydrogen bonding relationship of melting &
boiling point & solubility to intermolecular
forcesclassification of substances
UNIT 6:Liquids & Solutions
properties of solids phase changes & phase diagrams physical
properties of solutions concentration units colligative
properties
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give a brief overview of Thomsons experiments which led to the
discovery of the electron. (2.2) give a brief overview of
Rutherfords experiments which led to the discovery of the nucleus.
(2.2) describe the mass & charge of each of the 3 fundamental
subatomic particles: proton, electron & neutron. (2.2) explain
the terms: atomic number, mass number, molar mass, atomic mass
unit, isotope and atomic mass. (2.3, 3.2) describe the main
features of the Rutherford model of the hydrogen atom. (class
notes) define wavelength and frequency, and solve problems using
the relationship . (7.1) identify the regions of the
electromagnetic spectrum based on their frequencies and
wavelengths. (7.1) explain the difference between continuous and
line spectra. (7.3) describe the experimental setup to obtain the
atomic spectrum of hydrogen. (class notes, emission experiment)
give a brief outline of Plancks quantum theory and solve problems
based on Plancks equation . (7.1) describe the Bohr model of the
hydrogen atom. (7.3) use the Bohr equation to calculate the energy
of the electron in a given Bohr orbit. (7.3, emission experiment)
calculate E for the transition of an electron from one Bohr energy
level to another. Also calculate and associated with these
transitions. (7.3, emission experiment) recognize how calculations
of E, and for energy transitions in the hydrogen atom enabled Bohr
to explain the observed line spectrum. Also explain the importance
of the Bohr model in the eventual development of atomic theory.
(7.3, emission experiment) explain why, based on the Bohr model,
hydrogen gives a line, not a continuous spectrum. (class notes,
emission experiment) explain the shortcomings of the Rutherford
model of the hydrogen atom. (class notes) explain the fundamental
differences between the Bohr and Rutherford models of the hydrogen
atom. (class notes) recognize the difference between emission and
absorption spectra. (class notes)
Unit II: Atomic Theory(Chang, Ch. 2, 7 & 8)
Objectives:
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recognize the difference between the ground state & excited
states in the hydrogen atom spectrum. (class notes, emission
experiment) calculate the ionization energy of the hydrogen atom
(7.3, emission experiment) recognize that the Bohr model was only
successful for one electron species, and explain the general
limitations of the Bohr model. (7.3) outline the general features
of the quantum mechanical model of the hydrogen atom based on the
idea of the wavelike properties of matter. (7.5) outline the
contribution of each of the following scientists to the development
of the quantum mechanical model of the atom: Bohr (7.5), Heisenberg
(7.5), Schrdinger (7.5), De Broglie (7.4), Pauli (7.8). list the
quantum numbers: n, l, ml obtained by solving the Schrdinger wave
equation for the electron in the hydrogen atom. Predict the
allowable values of these quantum numbers. (7.6) explain the
meaning of wave function and orbital. (7.5) relate the shapes of
orbitals to the quantum number l. (7.6, 7.7) sketch the shapes of s
and p orbitals. (7.7) recognize the shapes of d orbitals. (7.7)
explain the meaning of electron spin and its relationship to the
quantum number ms (7.6) apply the Pauli exclusion principle and
Hunds rule to write the electron configuration and orbital filling
(box) diagrams for multielectron atoms. (7.8) explain and apply the
concept of effective nuclear charge. (7.9 & 8.3) identify the
s, p, d, and f blocks in the periodic table. (7.9) relate the
position of an element in the periodic table to its electron
configuration. (7.9) know the anomalous electron configurations for
first and second row transition metals. (7.9) explain and predict
the trends in atomic radii. (8.3) predict whether an atom is para
or diamagnetic based on its electron configuration. (7.8) give the
allowable sets of the four quantum numbers (n, l, ml, ms ) for each
electron in a given atom. (7.8)
Continuation
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2.1 Atomic Theory(Chang, 2.1)
Insight into the nature of matter:
Aristotle (384-322 BC)
Democritus (460-370 BC)
Aristotle believed that matter was infinitely divisible.
Democritus believed that matter was made up of tiny indivisible
particles that he called atoms
(from Greek, atomos, meaning indivisible).
Atoms are the basic unit of matter.
Atoms are extremely small (sheet of paper = 1 million atoms
thick, 1 drop of water = 1022 atoms)
Does matter have a basic/fundamental particle?
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2.1 Atomic Theory(Chang, 2.1)
J. Dalton (1766-1844)
John Dalton was an English scientist who created a hypothesis
about the nature of matter.
1. Elements are composed of extremely small particles called
atoms.
2. All atoms of a given element are identical (same size, mass
& chemical properties). The atoms of one element are different
from the atoms of all other element.
3. Compounds are composed of atoms of more than one element. In
any compound, the ratio of the #s of atoms of any 2 of the elements
is an integer.
4. A chemical reaction involves the separation, combination, or
rearrangement of atoms; no creation or destruction of atoms (based
on Lavoisiers Law of Conservation of Mass).
Note: there was no attempt to describe the structure of the
atom.
Daltons atomic model (1808) is based on the following
principles:
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2.2 Thomsons Experiment (Chang, 2.2)
J.J. Thomson, an English physicist , postulated the existence of
electrons based on his experiments with cathode ray tubes.
J.J. Thomson (1856-1940)
Using this experiment, Thomson determined
the charge to mass ratio of an electron:
e = -1.76 x 108 C/gm
The Plum Pudding Model
Thomson postulated that the atom was made of negatively charged
electrons in a positively charged cloud
(to maintain neutrality).
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2.3 The Millikan Experiment(Chang 2.2)
The American physicist, Robert A. Millikan, determined the
charge of an electron using the oil-drop experiment.
By having the charge suspended in air, he was able to determine
the charge on each oil drop. This led him to conclude that the
charge of 1 electron was:
-1.6022 x 10-19 C
This allowed the determination of the mass of a single
electron:
mass of an electron = charge = -1.6022 x 10-19 C = 9.10 x 10-28
gcharge to mass -1.76 x 108 C/g
R.A. Millikan (1868-1953)
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2.5 Rutherfords Experiment(Chang, 2.2)
The New Zealand physicist, Ernest Rutherford, bombarded alpha
particles (He2+, 7300 times the mass of an electron)through a gold
foil, expecting them to go straight through.
He proposed that the mass of the atom is concentrated in the
nucleus.
In separate experiments it was determined that the nucleus
contains positively charged particles called protons (equal but
opposite charge to the electron).
E. Rutherford (1871-1937)
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2.6 Summary of Subatomic Particles
The atom contains:
Nucleus, which consists of:
Electrons (same magnitude as proton but of negative charge).
Charge
Particle Mass (g) Coulomb Charge Unit
Electron 9.10938 x 10-28 -1.6022 x 10-19 -1
Proton 1.67262 x 10-24 +1.6022 x 10-19 +1
Neutron 1.67493 x 10-24 0 0
Protons (positively charge), Neutrons (almost same mass as
proton but no charge).
James Chadwick, a British physicist, discovered a new particle
in the nucleus of the atom, the neutron
(which comes from the latin neuter, which means neither one nor
the other), which has no charge
and holds the protons together.
J. Chadwick (1891-1974)
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2.7 Atomic Number & Mass Number(Chang, 2.3, Textbook
Problems: 2.14; 2.16; 2.18)
Atomic Number (Z): the # of protons in the nucleus(the # on the
periodic table).
XA
Z X is the symbol of the element.
The type of atom is determined by the # of protons (if an
element has a charge, it is due to a transfer of electrons and not
of protons).
In a neutral element:
In a cation: # of protons > # of electrons,(the element lost
electrons to become the ion).
In an anion: # of protons < # of electrons,
Mass Number (A): the # of protons plus the # of neutrons.
(the element gained electrons to become the ion).
# of electrons = # of protons.
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2.7.2 Exercises:
1. Fill in the missing information:
Element Symbol Z A # of
neutrons
# of
electrons
16 34
Iron Fe 56
17 16
80 120
Uranium U
2. Determine the # of protons, neutrons & electrons for each
of the following:
a) K
b) K+
c) Al3-
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2.8 Isotopes(Chang, 2.3)
Isotopes: substances with the same # of protons (hence, the same
element), but different # of neutrons (hence, different
masses).(from latin, iso = the same & topos = place: same place
in the periodic table)
Ex. H1
1
H1
2
H1
3
Hydrogen:(protium)
Deuterium:
Tritium:
1 proton, 1 electron, 0 neutron
1 proton, 1 electron, 1 neutron
1 proton, 1 electron, 2 neutrons
Note: Isotopes generally have the same chemical properties (same
# of protons & electrons), but have different physical
properties (different # of neutrons).
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2.9 Quantum Theory(Chang, 7.1, Textbook Problems: 7.4; 7.8;
7.10; 7.16; 7.18; 7.20; 7.22; 7.40; 7.42)
Properties of atoms & molecules are not governedby the same
physical laws as larger objects.
The German physicist Max Planck discovered that atoms &
molecules emit energy only in discrete energy called quanta (it was
previously believed that any amount of energy could be emitted or
released).
Waves
Wave: a vibrating disturbance by which energy is
transmitted.
Characterized by 3 factors:
Wavelength, (in nm): distance between identical points on a
successive wave.Frequency, ( in s-1 or Hz): # of waves that pass
through a particular point in 1 second.Amplitude: vertical distance
from the midline of a wave to the peak.
M. Planck(1858-1947)
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Continuation
It is also possible to describe a wave by its speed ():
= : wavelength (in nm) : frequency ( in s-1 or Hz)
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2.9.2 Electromagnetic Radiation
Electromagnetic radiation: describes how energy in the form of
radiation can be propagated through space as vibrating
electric& magnetic fields.
The speed of this waveis always the same& is equal to the
speed of light:
c = 2.9979 x 108 m/s
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2.9.1 Exercises:1. The wavelength of the green light from a
traffic signal is centered at 522 nm.
What is the frequency of this radiation?
2. True or False?
a) Blue light has a shorter wavelength than red light.
b) X-rays have lower frequencies than radio waves.
c) Microwaves have higher frequencies than gamma rays.
d) Visible radiation composes the major portion of the
electromagnetic spectrum.
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Photoelectric EffectPlanck gave the name quantum to the smallest
quantity of energy that can be
emitted (or absorbed) in the form of electromagnetic
radiation.
E = hn = h c
h is known as Plancks constant
(h = 6.626 x 10-34 Js)
Albert Einstein suggested that electromagnetic radiationbehaved
not only as waves but also as matter.
These particles of light are called photons and possess a
mass(theory of the dual nature of light, it can act as a wave &
as matter).
E = mc2Ephoton = h c
mphoton = E = hc/ = hc2 c2 c
A. Einstein(1879-1955)
Note: De Broglie postulated that matter can behave as waves as
well (c would have to be changed for which is the speed of the
object).
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2.9.2 Exercises:1. Sodium atoms have a characteristic yellow
color when excited in a flame. The color comes from theemission of
light of 589.0 nm. a)What is the frequency of this radiation? b)
What is the change inenergy associated with this photon? Per mol of
photons?
2. What is the wavelength of an electron (m = 9.11 x 10-31 kg)
travelling at 5.31 x 106 m/s? (1 J = 1 kg m2/s2)?
Conclusions: Matter & energy are not so distinct.
All matter shows particle properties as well as wave
properties.
Small matter (ex. photons) exhibit predominantly wave
properties.
Large matter (ex. baseball) exhibit predominantly particle
properties.
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2.10 Bohrs Theory(Chang, 7.3, Textbook Problems: 7.24; 7.30;
7.32; 7.34)
Emission SpectraWhen a substance is excited, either by heat or
an electrical current, it will emit light.
That is called the emission spectra of that substance.
There are 2 types of emission spectra:
Continuous spectrum:
Line spectrum:
Emission of all wavelength in the visible part of the
electromagnetic spectrum(white light is a continuous spectrum),
When passes through a prism it will look like a rainbow, the
colors will be separated.
Only specific wavelength are emitted,
Characteristic of an element (like the fingerprint of that
element).
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Line Spectra
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CVCV
CV
CV
orbit 4
orbit 3
orbit 2
orbit 1
nucleus
10.2 The Hydrogen Atom
The Danish physicist Neils Bohr imagined that the atom was like
a solar system with the electrons orbiting around the nucleus.
He imagined that electrons were only allowed to exist on certain
energy levels (i.e. the energy of the electron is quantized):
En = -RH 1n2
n is an integerRH equals 2.178 x 10
-18 J (Ryedbergs constant)
The most stable state is the ground state when n = 1.
All other states, n = 2, 3, 4 ... are known as excited
states(these are higher in energy than the excited state).
N. Bohr (1885-1962)( )
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When the hydrogen atom emits energy, the electron goes from a
higher energy level to a lower energy level (i.e. the E is
negative).
The radius of an orbit is proportional to n: the larger the n,
the larger the orbit, when the electron is farther from the
nucleus, it is higher in energy.
When the hydrogen atom gains energy, the electron goes from a
lowerenergy level to a higher energy level (i.e. the E is
positive).
10.2 The Hydrogen Atom
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The energy of the photons emitted is equal to the energy
difference between 2 energy levels:
The wavelengths emitted by hydrogen were separated in
different series based on their nfinal.
E = Efinal - EinitialE = (-RH) - (-RH)
n2f n2
i
E = -RH 1 - 1n2f n
2i
Continuation...
( )
The energy of an electron in another one electron species:
En = -RH Z2
n2(Z is the atomic #)
Note: these calculations only workfor 1 electron species.
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10.2.1 Exercises:1. Calculate the energy corresponding to the n
= 3 electronic state in the Bohr hydrogen atom.
2. For the excitation from the n = 1 to the n = 3 electronic
state in the hydrogen atom.a) Calculate the energy change.b) What
is the wavelength of the electromagnetic radiation associated with
this energy change?
3. Give examples of one electron species:
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10.2.1 Exercises:4. What would be the energy and wavelength (in
nm) of the light in the Lyman series when ni = 5?
5. Calculate the shortest and longest wavelength (in nm) in the
emission spectrum of the hydrogen atom for nf = 2 (Balmer
series).
6. Calculate the ionization of the hydrogen atom initially in
the ground state.
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2.11 Quantum Mechanical Model & Quantum Numbers(Chang, 7.5
& 7.6, Textbook Problems: 7.56; 7.58; 7.60; 7.62; 7.64;
7.66)
This model has a set of 4 quantum numbers (compared to 1 with
the Bohr model) to described the position of the electron.
Atoms emit or absorb energy when electrons move from one energy
state to another.
Electrons are found in orbitals (not orbits) which are 3
dimensional(these give the probability of finding the
electron).
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Principal Quantum Number (n):Can have a value from 1 to
infinity; it is always an integer.
It is the primary factor in determining the energy of an
electron.
It also determines the size of the orbital (average distance of
the electron from the nucleus).
The greater the n, the greater the energy & average
distance.
The primary quantum number is also referred to as a shell.
2
3
1
3
1
4
5
6
7
2
3
4
5
6
7
4
5
6
5
4
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Angular Momentum Quantum Number ():Can have a value from 0 n 1,
always whole numbers.
Determines the shape of the orbitals.
Is a secondary factor in determining the energy for
polyelectronic atoms.
The angular momentum quantum number is also referred to as
subshell.
Subshells are also referred to by letters according to their
#:
0 1 2 3 4 5
Name s p d F G h
(these letters refer to the shape)
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Magnetic Quantum Number (m):
Can have a value from - + .
Describes the orientation in space of the orbital.
The # of m possible is equal to 2 + 1.
S (l = 0)
d (l = 2)
p (l = 1)
f (l = 3)
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When = 0, s
When = 1, p
When = 2, d
Orbital Shapes:
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One last quantum number...
Superposing the Orbitals:
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Electron Spin Quantum Number (ms):
The value can be + or -.
Describes the spin of the electron.
Electrons act like tiny magnets and when then spin they create a
magnetic field.
There is 2 possible spinning motions for an electron (clockwise
& counter-clockwise).
S
d
p
f
-+
+ - Etc...
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2.11.2 Exercises:1. Determine the # of different orientations in
space for the s, p, d & f orbitals.
2. Which of the following sets of quantum numbers are allowed in
the hydrogen atom? For each incorrect set, state why it is
incorrect.a) n = 1, = 0, m = 1b) n = 2, = 2, m = 1c) n = 5, = 3, m
= 2d) n = 6, = -2, m = 2e) n = 6, = 2, m = -2
2.11.1 Exercise:
1. For principal quantum level n = 4, determine the # of allowed
subshells & give the name of each.
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2.11.3 Exercises:1. Which of the following sets of quantum #s
are not allowed? For each incorrect set, state why it is
incorrect.
a) n = 3, = 3, m = 0, ms = -
b) n = 4, = 3, m = 2, ms = -
c) n = 4, = 1, m = 1, ms = +
d) n = 2, = 1, m = -1, ms = -1
e) n = 5, = -4, m = 2, ms = +
f) n = 3, = 1, m = 2, ms = -
g) n = 3, = 2, m = -1, ms = 1
2. If each orbital can hold a maximum of two electrons (of
opposite spin) how many electrons can each of the following
subshells hold?
a) 2s
b) 5p
c) 4f
d) 3d
e) 4d
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2.12 Atomic Orbitals(Chang, 7.7, Textbook Problems: 7.68;
7.70)
The set of the 3 first quantum numbers describe the orbital.
The last quantum number describes the electron.
Energy of Orbitals
Orbitals that have the same energy are called degenerate.
For hydrogen:
All orbitals with the same value of n have
the same energy.
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Because of the penetration effect:
For polyelectronic atoms:
Energy of Orbitals
Ens < Enp < End < Enf
(i.e. the 2s orbital is lower in energy than the 2p orbital)
The penetration effect is caused by the other electrons which
are shielding the nucleus(the system must account for the
electron-electron repulsion).
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2.13 Electron Configuration(Chang, 7.8, Textbook Problems: 7.76;
7.78; 7.88; 7.90; 7.92; 7.96)
The electron configuration describes how the electrons are
distributedamong the various atomic orbitals.
1s1
It can also be described as an orbital box diagram:
In the ground state, the electron configuration of hydrogen
is:
1s1
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Aufbau Principle
As protons are added one by one to create the various elements,
so are electrons and these are placed in specific orbitals.
Aufbau principle: electrons are added to the lowest energy
orbitals first to create theground state of that element (from
German, meaning building-up).
The following chart describes which subshells are filled
first:
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Continuation
The other way to determine the order by which subshells are
filledis by following the periodic table:
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Pauli Exclusion Principle
The Austrian physicist Wolfgang Pauli established the principle
that no two electrons can have the same set of 4 quantum
numbers.
Since the first 3 quantum numbers describe the orbital and the
fourth quantum number can only have 2 values (+ & -):
there can only be 2 electrons per orbital, each of opposite
spin.
1. Write the electron configuration and box diagram for the
following atoms:
a) H
b) He
c) Li
d) Be
e) B
2.13.1 Exercise:
W. Pauli(1900-1958)
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Hunds Rule
The German physicist Friedrich Hermann Hund established the rule
wherein the lowest energy configuration is the one
having the maximum number of unpaired electron of parallel spin
in a particular set of degenerate orbitals.
1. Write the electron configuration and box diagram for the
following atoms:
a) C
b) N
c) O
d) F
e) Ne
2.13.2 Exercise:W. Hund
(1896-1997)
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Shorthand Notation
Shows in brackets the noble gas element that most nearly
precedes the element being considered:
Mg (Z= 12): 1s22s22p63s2 or [Ne]3s2
1. Write the shorthand notation for the following atoms:
a) Na
b) Al
c) Si
d) S
e) Ar
2.13.3 Exercises:
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2. Write the electron configuration and orbital box diagram for
the following atoms:
a) Sc
b) Ti
c) V
3. Write the electron configuration for the following atoms:
a) Mn
b) Fe
c) Co
d) Ni
e) Zn
f) Pb
g) Bi
2.13.3 Exercises:
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Exceptions:
There is a greater stability with a completely half-filled d
subshellor a completely filled d subshell.
To accomplish this, electrons can be taken from the previous s
orbitalto add to the d orbital.
Ex. Cr: 1s2 2s2 2p6 3s2 3p6 4s1 3d5
2p6
1s2
2s2
3s2
3p6
3d5
Cu: 1s2 2s2 2p6 3s2 3p6 4s1 3d10
4s1
2p6
1s2
2s2
3s2
3p6
3d10
4s1
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Valence ElectronsA group:
B group:
The valence electrons are those in the highest n(will be the
ones removes when a cation is formed).
The valence electrons are those in the highest n & those in
the (n 1) d orbital(those in the highest n will be removed first
when forming a cation).
Write the electron configuration for the following ions:
a) Na+
b) O2-
c) Ti2+
2.13.4 Exercise:
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2.15 Diamagnetism & Paramagnetism
Paramagnetic substances: contain unpaired electrons in
orbitals(these substances are attracted to a magnet).
Diamagnetic substances: all electrons are paired(these
substances are weakly repelled by a magnet).
1. Write the electron configuration of the following atoms and
determine if they are diamagnetic or paramagnetic:
a) O
b) Mg
c) N
d) Ar
e) Zn
f) Ti
2.15.1 Exercise:
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2.16 The Periodic Table
By 1869, a total of 63 elements had been discovered, and
scientists were studying their physico-chemical properties.
The chemist John Newlands classified these in order of
increasing atomic mass and noticed that their
properties repeated every 8 elements.
J. Newlands (1837-1898)
Newlands proposed the Law of octaves at the Chemical Society in
London, but was ridiculized because of the parellel he did with the
musical octave.
Contact: This isnt noise, this has structure!
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2.16 The Periodic Table
Drawing on Newlands discoveries, the Russian chemist Mendeleev
grouped in familiesall elements with similar properties &
created the 1st Periodic Table of Elements.
D. Mendeleev (1834-1907)
One advantages of his periodic table is that it classifies
elements both vertically(groups or families) and
horizontally (periods or rows).
Because of the limited knowledge, Mendeleev left gaps in the
periodic table of elements and correctly predicted the mass
and chemical properties of the yet unkown elements.
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Groups or families (vertical)
Periods or rows (horizontal)
Mass Number (A)
Relative Atomic Mass
The blue-print of all matter!
1A
2A
3B
8A
3A 7A6A5A4A
4B 5B 6B 7B 8B 1B 2B
2.16 The Periodic Table
A Group (Main Group): the representative elements (1A - 8A).
B group: the transition metals (1B-8B).
Inner transition metals: the 2 bottom rows (lanthanides &
actinides).
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2.16 The Periodic Table
Metals:
Metalloids:
Non-metals:
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Alkali metals
Alkali earth metals
Transition metals Noble gases
Halogens
Non metals
Other metals Lanthanides
Semi metals Actinides
2.16 The Periodic Table1A2A
3B
8A
3A 7A6A5A4A
4B 5B 6B 7B 8B 1B 2B
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2.17 End of Unit Exercises:1. The electron configuration of a
neutral atom in its ground state is [Xe]6s24f145d6.
a) What is the name and the symbol of the element?
b) How many unpaired electrons does the above element have?
c) Is the above element para or diamagnetic?
d) Is the above a metal or a non-metal?
2. For As (Z=33):
a) Write a complete set of quantum numbers for each of the
valence electrons.n m ms
b) Complete a similar table for the valence electrons of Fe
(Z=26),
c) Complete a similar table for the valence electrons of Cl
(Z=17),
d) Complete a similar table for the valence electrons of Cr
(Z=24),
e) Complete a similar table for all electrons of N (Z=7).
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3. Give the full electron configuration and box diagram for each
of the following species:
a) Mo (Z=42)
b) Cu (Z=29)
c) Zr (Z=40)
d) I-1 (Z=53)
e) Au (Z=79)
f) Zn2+ (Z=30)
g) Fe2+ (Z=26)
h) P-1 (Z=15)
i) (Z=81)
j) Fe3+ (Z=26)
k) Po2- (Z=84)
l) Cr (Z=24)
4. Determine if the species in 3) are paramagnetic or
diamagnetic.
End of unit exercises:
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5. Write the name and the symbol of one element which is
a/an:
a) Halogen in the 5th period,
b) First row transition metal,
c) Alkaline earth metal in the 2nd period,
d) Noble gas in the 3rd period,
e) Member of the 4a family,
f) Representative element of the 6th family,
g) Alkali metal in the 6th period,
h) Lanthanid,
i) Actinide,
j) Halogen with 74 neutrons.
6. Give the name & symbol of the element which corresponds
to each of the following information:
a) The heaviest halogen without d electrons,b) The first element
with five half-occupied 3d orbitals,c) The first element to have a
d electron,d) The first element to have a p electron,e) The first
element to have an f electron.
End of unit exercises:
