5 th Year Xmas Test Wed Dec 18 th 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 negativeCalled them electrons [after George Stoney of UCG]Magnet also deflects raysWorked out Charge/Mass ratioPlum Pudding model

Millikan Mass of Electron Oil Drop Experiment – charged oil drops with X-rays – measured charge to stop them falling – 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 NumberElement substance whose atoms all have same atomic number

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Bohr Niels Electrons orbit nucleus in shells – energy quantised – only certain values allowedMaximum number of electrons in shell = 2n2

Rutherford Called positive particles ProtonsSuspected other particles to cement nucleus together

Chadwick Showed Neutrons by bombarding Be with alpha particles

Sub-atomic ParticlesParticle Mass Charge Location

Proton 1 +1 NucleusNeutron 1 neutral NucleusElectron 0 -1 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

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IsotopesAtoms 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

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 M.Healy 2014/15

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 1sn = 2 2s, 2pn = 3 3s, 3p, 3dn = 4 4s, 4p, 4d, 4fn = 5 5s, 5p, 5d, 5f

Heisenberg’s Uncertainty PrincipleThe 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

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- Each d sublevel can hold up to 10 electrons

- Each f sublevel can hold up to 14 electrons

Electron configuration [pattern]

Aufbau PrincipleElectrons occupy the lowest available energy level

Order of filling

1s2s, 2p3s, 3p, 3d4s, 4p, 4d, 4f5s, 5p, 5d, 5f

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+ chargeEach 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

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p orbitals

Arrangement of electrons in orbitals of equal energy

Hund’s Rule of Maximum MultiplicityWhen two or more orbitals of equal energy are available, the electrons occupy them singly and then in pairs.

Final point is electrons spin in opposite direction

Pauli Exclusion PrincipleNo 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? M.Healy 2014/15

Na Mg Al Si P S Cl Ar

Na

Mg

Al

Si

PS

Cl

Ar

c

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

1st Ionisation energies

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. M.Healy 2014/15

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?

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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? ****************************

Chapter 4: Periodic Table

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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 simplerBoyles 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

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

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Modern Table v. Mendeleev’s

1. Arranged by atomic number2. More elements in modern – Noble Gases3. 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 – halogensGp 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

2311Na+

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.

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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

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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 substanceGroup of atoms chemically joined

IsotopesForms 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)

Atomic Number

Number of protons in an atom. Determines what the element is.

Mass number

Number of Protons + neutrons in an atom M.Healy 2014/15

Relative Atomic Mass

1. average of the mass numbers of the isotopes of the element.

2. as they occur naturally3. taking their abundances into account4. 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.

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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

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?

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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.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?

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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?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? M.Healy 2014/15

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 > 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

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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?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 of2 Charge3 Symbol4 Effect on atomic number5 Effect on Mass number

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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?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 M.Healy 2014/15

2. Find Mr of Mg (NO3) 2

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 10 23 particles of that substanceThis number 6 x 10 23 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

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One mole of oxygen molecules has a mass of 32gOne 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: 0.25moles5. 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 10 23

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

- What is the mass of one mole of sodium?... 23g

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- 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?..

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?

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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?

6 x 1023 = 1 mole1 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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