Download Relative Atomic Mass

5th Year Xmas Test Wed Dec 18th 2014
1. Atomic Theory:
Chapters 2,3,4,5,6,8,9.
Title -
2 EXPERIMENTS
Flame tests
Tests for anions in aqueous solutions:
Atomic Structure - History of Development
Democritus
Atoms = indivisible – 4 elements earth, wind, fire and water
Dalton John
Atomic Theory - element made of tiny particles – all particles of element identical –
compounds combinations of these particles
Crookes
Maltese Cross Experiment – Cathode Rays – radiometer
Thompson
Deflected cathode rays with charged plates – must be negative
Called them electrons [after George Stoney of UCG]
Magnet also deflects rays
Worked out Charge/Mass ratio
Plum Pudding model
Millikan
Mass of Electron
Oil Drop Experiment – charged oil drops with X-rays – measured charge to stop them fal
– used this with charge/mass ratio to work out mass
Rutherford
Gold Foil Experiment - Discovered nucleus – very small, dense, positively charged –
deflected alpha particles
Moseley
Each element has a characteristic positive charge – Atomic Number
Element substance whose atoms all have same atomic number
Bohr Niels
Electrons orbit nucleus in shells – energy quantised – only certain values allowed
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Maximum number of electrons in shell = 2n2
Rutherford
Called positive particles Protons
Suspected other particles to cement nucleus together
Chadwick
Showed Neutrons by bombarding Be with alpha particles
Sub-atomic Particles
Particle
Proton
Neutron
Electron
Mass
1
1
0
Charge
+1
neutral
-1
Location
Nucleus
Nucleus
Orbiting nucleus
Atomic Number = Number of Protons This determines what the element is.
Mass Number = Number of Protons plus the number of neutrons
Relative Atomic Mass is the average mass of the element as it occurs in nature [when the
isotopes and abundance are taken into account and compared to C12 being 12.
Number of Protons = Atomic Number
Number of Electrons = number of protons = atomic number [in atoms]
In ions add an electron for each minus charge and subtract an electron for each positive
charge
Number of Neutrons = Mass Number – Atomic Number
Isotopes
Atoms of an element which have the same atomic number but different mass numbers due to
different numbers of neutrons
Not all isotopes are radioactive but 14C is
Atomic Structure
Learn Bohr diagrams for first 20 elements
k = 2, l = 8, m = 8, n = >2
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Atomic spectra
3 types of spectra
Continuous – from sun
Line – Emission – from heated or shocked elements
– Absorption – frequencies removed as light passes through gaseous element
Emission Spectrum – how formed
-
Electron normally in lowest available energy level – Ground State – E1
Electron excited by heat or electricity [high voltage]
Jumps to higher energy level – Excited State – E2
Unstable
Drops back to lower level
Energy released as photon or Electro Magnetic Radiation
Frequency proportional to drop
(E2-E1) = hf
h is Planck’s Constant and f is frequency of radiation
lines formed are evidence of energy levels
Electron Configuration
n is the main energy level or Quantum Number
Number of electrons that a main energy level [shell] can hold is 2n2 [n = main energy
level number]
later found that main energy levels were divided into sublevels [up to 4 of these s, p, d and
f.
sublevels have different energies
n=1
n=2
n=3
n=4
n=5
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
5s, 5p, 5d, 5f
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Heisenberg’s Uncertainty Principle
The more accurately we determine the position of an electron the less
accurately we can determine its velocity.
Electrons have Wave-particulate duality
Have properties of both waves and particles at same time
Sublevels further divided into orbitals of equal energy
Each s sublevel contains 1 s orbital – spherical in shape
Each p sublevel contains 3 p orbitals - px, py, pz – each dumbbell shaped
and mutually at right angles
Each d sublevel contains 5 d orbitals
Each f sublevel contains 7 f orbitals
Orbital
Region in space around the nucleus of an atom in which electrons are
most likely to be found
Each orbital can hold up to 2 electrons
-
Each s sublevel can hold up to 2 electrons
-
Each p sublevel can hold up to 6 electrons
-
Each d sublevel can hold up to 10 electrons
-
Each f sublevel can hold up to 14 electrons
p orbitals
Electron configuration [pattern]
Aufbau Principle
Electrons occupy the lowest available energy level
Order of filling
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
5s, 5p, 5d, 5f
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Number of electrons [in an atom] = atomic number
Learn patterns for first 36 elements
e.g. Calcium
Atomic number 20 therefore 20 electrons
1s2, 2s2, 2p6, 3s2, 3p6,4s2
Exceptions
Chromium [4s1 3d5] and Copper [4s1 3d10]
Half full and full d sublevel is more stable
Electron configuration of ions
Each electron lost = 1+ charge
Each electron gained = 1- charge
Put pattern in square brackets with charge outside
e.g. Calcium ion Ca2+
Lost 2 electron thus 2+
[1s2, 2s2, 2p6, 3s2, 3p6]2+
Only patterns of ions of first 20 elements required
Arrangement of electrons in orbitals of equal energy
Hund’s Rule of Maximum Multiplicity
When two or more orbitals of equal energy are available, the electrons
occupy them singly and then in pairs.
Ar
Cl
P
S
Mg higher than general trend
due to full s sublevel and P is
higher than general trend due
to a half- filled p sub-level
Si
Mg
c
Al
Na
Na
Mg
Al
Si
P
S
Cl
Ar
1st Ionisation energies
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Final point is electrons spin in opposite direction
Pauli Exclusion Principle
No more than two electrons may occupy an orbital and to do so they must
have opposite spin
Each electron can be specified by 4 numbers
-
Principal Quantum Number
Sublevel
Orbital
Chapter 2-9
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?
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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:
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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?
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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.
Periodic Table - Atomic Theory Summary
Periodic Table
Boyle Robert
Earl of Meath – Elements can’t be broken down into anything simpler
Boyles Law – volume inversely proportional to pressure at constant temperature
Lavoisier Anton
Listed known elements
Law of conservation of Mass – matter neither created nor destroyed in a chemical
reaction simply rearranged.
Davy Humphrey
Discovered new elements Na, K Ca etc using Electrolysis
Dalton John
Atomic Theory - element made of tiny particles – all particles of element identical –
compounds combinations of these particles
Dobereiner
Triads – groups of 3 elements – middle one intermediate
Newlands
Law of Octaves – properties repeat every eighth elements when arranged in order of
mass – no noble gases
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Mendeleev Dmitri
Periodic Table – Arranged by property space rather than weight – left gaps – predicted
properties of elements to go in gaps – eka silicon (germanium)
Moseley Henri
Atomic number – number of protons
Modern Table
v.
Mendeleev’s
1. Arranged by atomic number
2. More elements in modern – Noble Gases
3. No gaps – Mendeleev left gaps to make elements fit into proper column – In
a few cases reversed the order of elements so they fitted into groups with
similar properties
4. Transition elements in a separate block
Groups
Vertical columns, Gp I – alkali metals; Gp II – alkaline earth metals , Gp VII –
halogens
Gp 0 – Noble gases – same number of electrons in out shell – similar properties
Periods
Rows – outer shell filling
Blocks
s, p, d [transition elements] , f
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?
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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?
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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 x 1023)
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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
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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.
Chapter 5: Bonding & Shapes of molecules
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1. What is meant by the octet rule?
2. Who were the 2 American chemists that proposed this
rule to predict how elements combine together?
3. Give 2 exceptions to the octet rule.
4. What is an ion?
5. What is an ionic bond?
6. What is the formula for calcium fluoride? Lithium
sulphide?
7. Give the formulae for the following complex ions:
carbonate, ammonium, sulphate.
8. Why do transition elements have variable valency?
9. What does –ate mean at the end of a formula?
10.
What is the definition of a transition metal?
11.
Why are scandium and zinc not considered true
transition elements?
12.
What is a molecule?
13.
What is a ‘lone pair’?
14.
What element is the standard by which valency is
measured?
15.
What is meant by the term valency?
16.
What is a sigma/pi bond?
17.
List 3 properties of a covalent compound.
18.
What is recrystallisation?
19.
Give two benefits of knowing a compound’s melting
point.
20.
What does VSEPR mean?
21.
Define electronegativity.
22.
Who set up the electronegativity scale?
23.
If electronegativity values are between 0 and 0.4
what does this tell you about the type of bonding?
24.
Name 2 polar substances.
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25.
What is the shape of the ammonia molecule?
26.
Name a non-metal that exists as single atoms.
27.
Which 2 of the first 36 elements would form the
compound with the greatest ionic character?
28.
What type of bond exists between the molecules
of hydrogen chloride.
List the 3 types of intermolecular forces
Chapter 5: Bonding & Shapes of molecules
29.
What is meant by the octet rule?
30.
Who were the 2 American chemists that proposed
this rule to predict how elements combine together?
31.
Give 2 exceptions to the octet rule.
32.
What is an ion?
33.
What is an ionic bond?
34.
What is the formula for calcium fluoride? Lithium
sulphide?
35.
Give the formulae for the following complex ions:
carbonate, ammonium, sulphate.
36.
Why do transition elements have variable valency?
37.
What does –ate mean at the end of a formula?
38.
What is the definition of a transition metal?
39.
Why are scandium and zinc not considered true
transition elements?
40.
What is a molecule?
41.
What is a ‘lone pair’?
42.
What element is the standard by which valency is
measured?
43.
What is meant by the term valency?
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44.
What is a sigma/pi bond?
45.
List 3 properties of a covalent compound.
46.
What is recrystallisation?
47.
Give two benefits of knowing a compound’s melting
point.
48.
What does VSEPR mean?
49.
Define electronegativity.
50.
Who set up the electronegativity scale?
51.
If electronegativity values are between 0 and 0.4
what does this tell you about the type of bonding?
52.
Name 2 polar substances.
53.
What is the shape of the ammonia molecule?
54.
Name a non-metal that exists as single atoms.
55.
Which 2 of the first 36 elements would form the
compound with the greatest ionic character?
56.
What type of bond exists between the molecules
of hydrogen chloride.
List the 3 types of intermolecular forces
Electronegativity
The relative attraction that an atom in a molecule has
for the shared pair of electrons in a covalent bond.
Invented by Linus Pauling
Difference in electronegativity and type of bonding
= 0 – 0.4
Pure Covalent bond
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> 0.4 < 1.7 Polar Covalent
> 1.7
Ionic
Chapter 6: Anion Tests
Identifying Anions [negative ions]

SO42-, Add Barium chloride / white precipitate / of barium sulphate / doesn’t
dissolve in dil. HCl

SO32-, Add Barium chloride nitrate / white precipitate / of barium sulphite /
does dissolve in dil. HCl

CO32-, gives CO2 when dil. HCl added / CO2 turns lime water milky / gives
white precipitate / of Magnesium carbonate / when magnesium sulphate
added

HCO31-, gives CO2 when dil. HCl added / CO2 turns lime water milky / does
not give white precipitate / of Magnesium carbonate / when magnesium
sulphate added

NO31-, add Iron(II)sulphate solution / Add conc. H2SO4 / at angle down test
tube / acid sinks to bottom / brown ring at interface is a positive result

PO43-, Add conc. nitric acid to ammonium molybdate / add mixture to nitrate /
heat gently in water bath / yellow precipitate is positive result

Cl1- Add acidified silver nitrate solution / white precipitate / of silver chloride /
re-dissolves in / ammonia solution
1. What is the Law of Conservation of Mass?
2. Be able to balance simple equation.
3. What is an anion?
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4. What is the name given to a material that settles out
of solution?
5. Give reagents/chemicals needed to test for all 7
anions. (Usually 2)
6. Give equations where necessary.
Radioactivity
Chapter 8
Radioactivity – There are three types of
radioactivity: alpha, beta and gamma. Be able to write
brief notes on each type of radiation using the
following guidelines.
1 What each type consists of
2 Charge
3 Symbol
4 Effect on atomic number
5 Effect on Mass number
6 Penetrability
1. Who discovered radioactivity?
2. Who discovered polonium and radium?
3. What is radioactivity?
4. What are the three types of radiation?
5. What instrument is used to detect radiation?
6. What units is radioactivity measured in?
7. What is half-life?
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8. What is an alpha particle? Beta particle?
9. How can they be stopped?
10.
Give a source of each of the three particles.
11.
What is a radioisotope?
12.
Complete nuclear equations to show the
action of alpha and beta particles.
Chapter 9: The Mole
Relative Molecular Mass of a
substance is the mass of one molecule
of that substance compared with one
twelfth of the mass of the carbon-12
isotope
1. Find Mr of CuSO4
Answer: 159.5
2. Find Mr of Mg (NO3) 2
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Answer: 148g
A Mole
is a measure of the amount of a
substance
One mole of a substance is the amount
of that substance that contains 6 x
1023 particles of that substance
This number 6 x 1023 is known as the
Avogadro Constant
A mole of a substance is either its
relative atomic mass or its relative
molecular mass (which ever is
suitable) expressed in grammes.
One mole of carbon has a mass of 12g
One mole of oxygen molecules has a
mass of 32g
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One mole of CuSO4 has a mass of
159.5g
No. of moles = m/M
= mass in grammes/
mass of one mole
3. How many moles are in 16.6g of
CO2
Answer: 0.377 moles
4. How many moles are in 10g of
NaOH?
Answer: 0.25moles
5. What is the mass of 0.05 moles
of CuSO4
Answer: 0.05 x 159.5 = 7.9moles
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6. If 0.2 mole of a compound has a
mass of 9.2g, what is the relative
molecular mass of the compound?
Answer 9.2 /0.2 = 46g
No. of Particles
One mole of any substance contains
6 x 1023
7. How many (a) molecules (b) atoms
are there in 42.4g of Na2CO3
Chapter 9: The Mole
1. Give a definition of a mole.
2. What is the Relative Molecular Mass of a
compound?
3. What is Avogadro’s number?
4. Calculations
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- What is the mass of one mole of sodium?...
23g
- What is the mass of a molecule of oxygen?
32g
- What is the Relative molecular mass of
water?
- .. 18g
- How many moles are in 10g of sodium?
- ..10/23 = .43moles
- How many atoms are in 10g of sodium?..
M.Healy
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10/23 = .43 x 6 x 1023
= 2.5 x 1023
5. What mass of chromium has the same number
of atoms as 8g of calcium?
8/40 = 0.2moles 1 mole Cr = 52g 0.2moles Cr =
10.4
6. How many atoms are present in 0.12g carbon.
0.12/12 = 0.01moles
1 mole carbon = 6 x 1023
0.01 x 6 x 1023
7. How many moles are in 2.6 x 1012 atoms of
sodium?
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6 x 1023 = 1 mole
1 atom of sodium = 1/6 x 1023
2.6 x 1012 = 2.6 x 1012
________
6 x 1023
= 0.43 x 1012-23
= 4.3 x
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