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Chemistry ATAR Year 11 > 1 - Properties and Structure of Atoms > Flashcards

1 - Properties and Structure of Atoms Flashcards

1
Q

State the History of the Atomic Theories and Models, Indicating the Model Name and Inventor. Give the Year of Discovery if Possible.

A
  1. Solid Sphere Model (1803) by John Dalton
  2. Plum Pudding Model (1904) by J. J. Thompson
  3. Nuclear Model (1911) by Ernest Rutherford
  4. Planetary (Bohr’s) Model (1913) by Niels Bohr
  5. Quantum Model (1926) by Erwin Schrödinger
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2
Q

What Must be Shown in a Labelled Diagram of an Atom?

A

Element, Number of Protons and Neutrons, Organisation of Electrons (dots) in Orbitals/Energy Shells around a Nucleus.

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3
Q

Describe the Proton, Its Symbol, Location in the Atom, Relative Charge and Relative Mass

A

Proton (p) is a positively charged subatomic particle.
Found in the Nucleus
Charge of 1+
Relative Mass of 1

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4
Q

Describe the Neutron, Its Symbol, Location in the Atom, Relative Charge and Relative Mass

A

A Neutron (n) is a neutral subatomic particle found.
Found in the Nucleus of an Atom
Relative Charge of 0
Relative Mass of 1

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5
Q

Describe the Electron, Its Symbol, Location in the Atom, Relative Charge and Relative Mass

A

Electron (e) is a negatively charged subatomic particle.
It is located in orbitals and energy shells around the nucleus
Relative Charge of 1-
Relative Mass of 1/1840

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6
Q

How are Elements Represented by X, A and Z

A

X = Atomic Symbol (Name Abbreviation)
A = Mass Number (p+n, electrons are so tiny their masses are negligible)
Z = Atomic Number (p)

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7
Q

What Holds an Atom Together?

A

The Electrostatic Force of Attraction between the Positively Charged Nucleus and Electrons Around the Nucleus.

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8
Q

What Holds a Nucleus Together?

A
  1. Electrostatic Force of Repulsion between Protons (Same Charge); pushing them away, but into each other (consider a bunch of protons)
  2. Strong Nuclear Force - Short Ranged Force between all particles, regardless of charge.

Balance of the Above Forces Keeps the Nucleus Stable.

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9
Q

What are Isotopes?

A

Isotopes are Atoms of an Element with the Same Number of Protons but Different Number of Neutrons.
Such that they have different Mases

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10
Q

How can Isotopes be Represented?

A

(superscript) A X (IUPAC)

Or

X - A

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11
Q

What are Physical Properties?

A

Observable or Measurable Features
E.g. Density, Solubility, Colour, Mass

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12
Q

What are Chemical Properties?

A

The Properties that Determine how a Substance Reacts in a Chemical Reaction

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13
Q

Discuss and Explain the Difference in Chemical and Physical Properties of Isotopes of an Element.

A

Isotopes will have:
- Different Physical Properties due to their different masses (different number of neutrons)
- Similar Chemical Properties as the electron configuration of isotopes of an element remain constant (and the electron arrangement determines the chemical properties of an element, particularly the valence electrons)

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14
Q

What is Relative Atomic Mass?

A

Relative Atomic Mass [or Ar (subscript)], compares the mass of an atom to 1/12 of the mass of C-12 as actual atomic masses are extremely small.

There is No Units!

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15
Q

How is Relative Atomic Mass Calculated?

A
  1. Refer to the Periodic Table of Elements
  2. Use the Formula:

Ar = (Abundance % * Mass Number) + (Abundance % * Mass Number) + …

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16
Q

What is Relative Molecular Mass and Relative Formula Mass? What is the Difference?

A

Relative Molecular Mass (Mr) compares the mass of a molecular compound to 1/12 of a C-12 atom. It is calculated by adding the appropriate relative atomic masses of each element.

Relative Formula Mass (Mr) compares the mass of an ionic substance to 1/12 of a C-12 atom. It is calculated by adding the appropriate relative atomic masses of each element in its simplest ratio.

Mr(__) = Ar(__) + Ar(__) + …

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17
Q

Given a Naturally Occurring Piece of Magnesium contains a mixture of isotopes: Mg-24, Mg-25, Mg-26 and the relative atomic mass of Mg is 24.3, which isotope is the most abundant?

A

Mg-24 is the most abundant isotope in the piece of magnesium as it is closest to the rounded value of Magnesium’s relative atomic mass. This is because the relative atomic mass of an element reflects its isotopic composition and abundances.

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18
Q

Example:

Find the Relative Atomic Mass of Carbon given the Relative Abundance of:
C-12 is 98.9%
C-13 is 1.1%

A

Ar (C) = (0.989 * 12) + (0.011 * 13)
= 12.01

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19
Q

Explain how you would calculate the Relative Atomic Mass of 2 unknown isotopes of 3, given the Relative Atomic Mass of the Element?

A

You would assign one unknown isotope to be x%, and the other to be 100% - x% - known isotope % and equate to solve.

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20
Q

Example:

Calculate the Relative Molecular Mass of Dinitrogen Tetroxide

A

Mr(N2O4) = 2 * Ar(N) + 4 * Ar(O)
= 2(14.01) + 4(16.00)
= 92.02

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21
Q

Example:

Calculate the Relative Formula Mass of Ammonium Oxide

A

Mr[(NH4)2O] = 2Ar(N) + 8(H) + 1*Ar(O)
= 2(14.01) + 8(1.008) + (16.00)
= 52.084

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22
Q

What is the Max Number of Electrons that Can fit in Each Shell? What is the General Rule used for this?

A

2n^2, where n=number of energy shells
When n=
1: 2
2: 8
3: 18
4: 32

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23
Q

How do Electrons fill Energy Shells?

A

Electrons always fill the Energy Shell that Requires the least amount of energy to occupy.

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24
Q

What is an Electron Configuration and how is it structured?

A

The Electron Configuration shows the arrangement of electrons around the nucleus of an atom.

For example, Calcium [2, 8, 8, 2]

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25
Q

Explain Why the Energy Shells do not always fill up Entirely before Electrons begin to fill in the next shell?

A

Electrons will always fill up the energy level that requires the least amount of energy. However, an anomaly is seen when electrons begin to fill up the 3rd energy shell to 8, before filling up the 4th energy shell to 2, then filling the 3rd energy shell again. This can be explained by the sub-shells (regions of energy levels containing orbitals of the same energy) and orbitals. Essentially, The Third Shell Fills to 8 Electrons, then the Fourth Shell Fills to 2, then the Third Shells Fills until 18. This is entirely due to energy usage to bond electrons to an atom.

26
Q

Who Developed the Period Table of Elements and When?

A

Dimitri Mendeleev in 1869

27
Q

Explain the Terms: Groups and Periods in reference to the Periodic Table of Elements.

A

Periods = Horizontal Rows
Groups = Vertical Rows

28
Q

How are Elements Arranged on the Periodic Table?

A

Elements are arranged in increasing order of atomic number (number of protons) and and grouped based on similar properties (along same group due to same electron configurations)

29
Q

Give Common Names for the Groups of the Periodic Table of Elements and Properties

A

Group 1: Alkali Metals - Highly Reactive, Readily with Halogens
Group 2: Alkaline Earth Metals - Shiny, Silvery, White Metals Found Commonly on Earth; somewhat reactive
Group 3-12: Transition Metals - Metals of Varying Properties
Group 17: Halogens - Highly Reactive; Lethal and Toxic
Group 18: Noble Gases - Odourless, Colourless, Low Reactivity

30
Q

What is a Metalloid?

A

An element sharing properties of metals and non-metals.

31
Q

What is Core Charge? How is it Calculated and What is its Trend?

A

The Measure of Attraction between the valence electrons and nucleus of an atom.
Calculated: p - inner shell e

Relationship:
- Constant Down a Group (indicated by the units digit of the group number) - Proportion of Electrons and Nucleus Size increasing is the same.
- Increases Across a Period (Electrons are closer to the nucleus and more valence electrons)

Lower Core Charge = Higher Reactivity

32
Q

Describe and Explain the Periodic Table Trend of:

Valence Electrons

A

Outermost Electrons

  • Increases Right of a Period
  • Constant Down a Group
33
Q

Describe and Explain the Periodic Table Trend of:

Number of Energy Levels

A

Number of Shells Electrons can Fill

  • Remains Constant Across a Group
  • Increases Down a Group
34
Q

Describe and Explain the Periodic Table Trend of:

Atomic Radius

(and Ionic Radius)

A

Atomic Radius - Distance from Nucleus to Outermost Electron of an atom

  • Increases Left of a Period (as when core charge increases, electrons become more closely attracted to the nucleus)
  • Increases Down a Group (more energy levels filled, so electrons are further from the nucleus).

Ionic Radius - Distance from nucleus to outermost electrons of an ion

  • Same as atomic radius, as proportionally, the distance remains constant
  • Cations get smaller (reduced electrons)
  • Anions get bigger (increased electrons)
35
Q

Describe and Explain the Periodic Table Trend of:

First Ionisation Energy

A

Amount of Energy Required to Remove One Electron from an atom in Gas Phase

  • Increases Right of a Period (Core Charge increases and more energy is required to overcome the electrostatic attraction between electrons and the nucleus)
  • Decreases Down a Period (as electrons are further away from the nucleus and require less energy to remove)
36
Q

Describe and Explain the Periodic Table Trend of:

Electronegativity

A

Tendency of an Atom to Attract a Bonding Pair of Electrons.

  • Increases Right of a Period (Core Charge Increases, so less energy is required as electrons are closer to the nucleus and attraction is stronger)
  • Decreases Down a Group (Electrons are further away from the nucleus due to increasing energy shells, so electrons are less attracted to the nucleus

Excludes; Noble Gases

37
Q

Describe and Explain the Periodic Table Trend of:

Electron Affinity

A

The Ability of an Atom in Gaseous State to Accept an Electron and Form an Anion

  • Increases Right of a Period
  • Decreases Down a Group

Because an atom is able to accept an electron more willingly is its core charge is greater, with more protons and electrons closer to the nucleus (smaller atomic radius)

38
Q

Describe and Explain the Periodic Table Trend of:

Metallic Character

A

The Measure of how much an atom behaves like a metal
Such as; shiny, brittle, malleable, etc.

Particularly, Metallic Character is Evident when the Attraction between Valence Electrons and Nucleus is Weaker, Enabling a Loss of Electrons

  • Increases Left of a Period (lower core charge, valence electron is less likely to be stable)
  • Increases Down a Group (Atomic Radius Increases as electrons are further away from the nucleus and are less attracted)
39
Q

What is ‘Ground State’

A

Describing Electrons as being in the lowest possible energy levels in an atom

40
Q

What is an ‘Excited State’?

A

When the electrons in an atom are at a higher energy shell than normal (not the ground state)

41
Q

Explain what happens to an atom when energy is applied to it (i.e. heat)

A

An atom is usually in ground state, such that its electrons are at the lowest energy level possible. When energy is applied to an atom in ground state, the electrons gain energy and move up energy levels to an excited state. However, this state is unstable so electrons return back to their original state and release energy in the form of light.

42
Q

Why do different elements emit different coloured light when heated or energy is added?

A

All elements have unique sets of energy levels and electron arrangements. Hence, the absorption and emission of energy is consistent for all atoms of the same element and release unique wavelengths of light when returning from an excited state to a ground state.

43
Q

When Heated, what Light is Emitted from:

Calcium

A

Orange

44
Q

When Heated, what Light is Emitted from:

Cobalt (II)

A

White Flashes

45
Q

When Heated, what Light is Emitted from:

Copper (II)

A

Bright, Fluro Green

46
Q

When Heated, what Light is Emitted from:

Barium

A

Pale Blue-Green

47
Q

When Heated, what Light is Emitted from:

Magnesium

A

Bright White

48
Q

When Heated, what Light is Emitted from:

Nickel

A

Silver White

49
Q

When Heated, what Light is Emitted from:

Potassium

A

Lilac (Pink)

50
Q

When Heated, what Light is Emitted from:

Sodium

A

Strong, Persistent, Solid Yellow

51
Q

When Heated, what Light is Emitted from:

Strontium

A

Red

52
Q

What is the Difference Between Spectroscopy and Spectrometry?

A

Spectroscopy - Study of radiated energy and the electromagnetic spectrum; and how they interact

Spectrometry - Use of Spectroscopy so quantifiable measurements can be made.

53
Q

What is the Difference Between Absorption and Emission?

A

Absorption - intake of electromagnetic spectrum (as light) as energy

Emission - release of electromagnetic spectrum (as visible light) as energy

54
Q

Describe the Method of Atomic Emission Spectroscopy (AES)

A

AES can be demonstrated by the flame test, or using an AES spectroscope.

Method:
1. A substances is heated and light is released.
2. The light is passed through a lens and prism to separate the wavelengths of light
3. The wavelengths are recorded into a photographic plate

55
Q

Describe the Results and Use of the Atomic Emission Spectroscopy (AES)

A

The AES Machine is used to produce an Emission Spectrum; which shows the wavelengths of light produced on a black photographic plate.

It is a qualitative measurement and only identifies which element is being tested.

56
Q

Describe the Method of Atomic Absorption Spectroscopy (AAS).

A

AAS is used to test the amount of specific wavelengths of light absorbed.

  1. A hollow cathode lamp containing the metal being analysed emits the specific wavelength of light particular to that element.
  2. The sample being tested is vaporised and its atoms turn into ions in the flame. These ions in an excited state will absorb the specific frequency wavelengths produced by the hollow cathode lamp and the remaining light will be emitted.
  3. The light is passed through a prism/monochromator that isolates the specific wavelength of light being tested.
  4. The degree of light absorption is measured by a photomultiplier
57
Q

Describe the uses and function of the Atomic Absorption Spectroscopy (AAS)

A

AAS is used as a qualitative tool as it can identify the presence of a particular metal atom by inverting the Absorption Spectrum produced.

It is also used as a quantitative tool as the degree of light absorption is proportional to the concentration of metal in the sample; indicating an increase in the light absorb means there are more atoms present in that sample.

Hence, a calibration graph comparing absorption value (y-axis) and Metal Ion Concentration (x-axis) can be produced.

58
Q

Describe how you would use a Calibration Graph from Atomic Absorption Spectroscopy.

A

If given an absorbable value, you can figure out the corresponding (expected) concentration of the metal ion in the sample.

59
Q

Describe the Method for Mass Spectroscopy.

A

Prior - the Sample being tested is vaporised into gas phase molecules.

  1. Ionisation - the vaporised sample is bombarded with high speed electrons or ultraviolet light. This produces positive ions as electrons are knocked off.
  2. Acceleration - the positive ions are accelerated through an electric field to move at high-velocities.
  3. Deflection - the ions pass through a magnetic field and undergo deflection according to the mass-to-charge ratio (m/z). Lighter ions are deflected more than heavier ions.
  4. Detection - the ions are projected onto a sensitive plate and detects the mass and number of times ions strikes a certain location; recording the data into a mass spectrum.
60
Q

Describe the uses and function of Atomic Absorption Spectroscopy (AAS).
How is the produced graph read after AAS is used on a specific sample?

A

AAS is used to determine the abundance of isotopes (isotopic composition).

This is achieved by the detection plate producing a mass spectrum graph that compares the percentage abundance to mass-to-charge ratio. The spikes in the graph can indicate the most abundant isotopes, the percentage abundance of each isotope, and can be further extrapolated to measure its average atomic mass.

Chemistry ATAR Year 11 (2 decks)

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