Chapter 2
Atomic Theory
A brief history of the atom also with
special attention to Dalton, Crookes,
Thompson, Millikan, Rutherford and
Chadwick
Q.1 What is meant by diffusion? Give e.g.
Q.2 What is formed when hydrochloric acid and
ammonia react?
Q.3 Who was the first man to put forward an atomic
theory
Q.4 What material did Rutherford hit with alpha
particles in his famous experiment that led to the
discovery of the nucleus?
Q.5 What is the mass of a neutron?
Q.6 What was Dalton’s theory?
Q.7 Who studied the discharge of electricity
through gases?
Q.8 What are cathode rays?
Q.9 Draw a cathode ray tube?
Q.10 What name is given to the negative/positive end
of a battery?
Q.11 Who discovered the electron?
Q.12 What does e/m mean?
Q.13 Who devised an expt. to measure the charge on
the electron?
Q.14 What did Thomson’s model of the atom look
like?
Q.15 Describe Rutherford’s expt. Give his findings.
Q.16 Who discovered the proton?
Q.17 Who discovered the neutron?
Q.18 Describe the expt.that led to the discovery of
the neutron?
Q.19 What is the mass of the proton/the electron?
Q.20 What did Thomson discover about cathode
rays?
Q.21 What contribution did Stoney make in the
discovery of the electron?
Q.22 Compare the mass, charge and location of the
proton, neutron and electron.
Chapter 3:
Arrangement of Electrons in an atom
1. Who was the first scientist to give information about
the arrangement of electrons in an atom?
2. What is a continuous spectrum/line spectrum?
3. Name two ways in which spectra can be seen.
4. Describe an expt. to investigate the flame colours of
different salts.
5. Name two parts of a spectrometer.
6. What colour is emitted by lithium, potassium, barium,
strontium, copper and sodium?
7. What salt commonly causes contamination?
8. How do you reduce contamination?
9. What is an energy level?
10.
What is a quantum of energy?
11.
What is Heisenberg’s Uncertainty Principle?
12.
Who worked out the likely probability of finding a
particular electron in an atom? (mathematically)
13.
What is the shape of an s/p orbital?
14.
How many type of p orbital exist? Draw, give
letters to label each.
15.
What is the ground state of an atom?
16.
How do you work out how much energy is emitted
when an electron falls back down to the ground state,
from its excited state.
17.
Give detailed description (using diagrams) of how
elements are able to produce their own particular line
spectra.
18.
What is the Balmer series?
19.
What is the definition of an orbital?
20.
What did De Broglie say about moving particles?
21.
What is an absorption spectrum? Give some of its
uses (two)
22.
What does LASER stand for?
23.
Give another example of a piece of equipment that
makes use of electron transitions.
What does ‘h’ stand for in the equation relating energy of
light and its frequency.
24. Name the instrument used to study emission line
spectra?
25. Give another word to explain ‘quantised’, with reference
to energy of an electron?
26. E = hf. What does ‘f’ stand for?
27. What electron transitions give rise to lines in the visible
spectrum?
28. What is the maximum number of electrons that can
occupy energy Level 3?
29 Which of the following orbitals is spherical: ‘s’, ‘p’, ‘d’ or
‘f’
30. Give the name of two ‘heavy’ metals that might be
found by Atomic Absorption Spectrometer in water
analysis?
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Chapter 4: Periodic Table
The Periodic Table – Know the contributions made
by Dobereiner (triads), Newlands (octaves),
Mendeleev and Mosely..
Know the difference between Atomic number and
Relative Atomic Mass, group and period, metals
and non-metals and gases. Be familiar with the
following groups – alkali metals, alkaline earth
metals, halogens and noble gases.
Calculation of the relative atomic mass from the
percentage isotopes should be practised.
1. Give explanation and example of how Dobereiner
and Newlands grouped elements.
2. Why would Newland’s classification not work
today? Give the main reason.
3. In what order did Mendeleev arrange elements?
4. Who changed this order?
5. Give three differences between Mendeleev’s
table and the modern day one?
6. What is meant by ‘eka silicon’?
7. Who discovered sodium and potassium?
8. What is the Periodic Law?
9. How many (a) electrons (b) neutrons are in
23
+
11Na
10.
What is Relative Atomic Mass?
11.
What is a Mass Spectrometer used for?
(Give 2)
12.
Who built the first Mass Spectrometer?
13.
Give the Principles behind the operation of
the Mass Spectrometer.
14.
Name the parts and be able to draw the
instrument.
15.
What is an isotope? Give an e.g.
16.
Calculate the relative atomic mass of an
element give its mass and % abundance.
17.
Write the electronic configuration of an
element..
18.
Write the s,p configuration of an element..
19.
What is the order in which the sub levels are
filled?
20.
What does isoelectronic mean?
21.
What is Pauli’s Exclusion Principle?
22.
What is Hund’s Rule?
23.
What is the Aufbau Principle?
24.
How many electrons can be accommodated in
the p orbitals/d orbitals?
25.
Why is the arrangement of electrons in
potassium 2,8,8,1 and not 2,8,9,?
26.
Write the s,p, configuration for Copper and
Chromium? What is different about these?
DEFINITIONS Chapters 2,3,4.
Atom
Smallest particle of matter that can exist by itself
Matter
Anything that occupies space
Element
Substance made up of one type of atom –
can’t be broken into anything simpler by chemical
means
Molecule
Smallest particle of substance that shows properties
of that substance
Group of atoms chemically joined
Isotopes
Forms of element with different mass number due to
different numbers of neutrons
The Mole
1 mole = 1 mole
= RMM (relative molecular mass) in grams
= Avogadro’s number (6*1023)
= 22.4L of any gas at STP [ 273 K and 760mm
of Hg ]
Avogadro’s Law
1 mole of any gas at STP occupies 22.4 L
Atomic Number
Number of protons in an atom. Determines what
the element is.
Mass number
Number of Protons + neutrons in an atom
Relative Atomic Mass
1. average of the mass numbers of the isotopes of the
element.
2. as they occur naturally
3. taking their abundances into account
4. expressed on a scale on which atoms of the carbon
12 isotope have a mass of exactly 12 units.
Relative Molecular Mass
1. The sum of the relative atomic masses of all the
atoms in a molecule of the compound.
2. The mass of one molecule of that compound
compared with one twelfth of the mass of the
carbon 12 isotope.
3. Mass of one mole of a compound = Relative
Molecular Mass in grams.
Energy Level
The fixed energy value that an electron in an atom
may have.
Atomic Orbital
The region in space within which there is a high
probability of finding an electron.
Hund’s Rule of Maximum Multiplicity
When 2 or more orbitals of equal energy are
available, the electrons occupy them singularly
before filling them in pairs.
Aufbau Principle
Electrons occupy the lowest available energy level.
Pauli Exclusion Principle
No more than 2 electrons may occupy an orbital and
they must have opposite spin.
Heisenberg’s Uncertainty Principle
The more accurately we know the position of a
particle the less accurately we know its velocity.