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Chemistry > Chapter Three > Flashcards

Chapter Three Flashcards

Electrons and the periodic table

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1
Q

How do electrons orbit the nucleus

A

In fixed circular orbits which correspond to specific energy shells/levels of the atom

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2
Q

What do electron orbits of a larger radii mean

A

They correspond to a higher energy level

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3
Q

Can electrons occupy the space between energy levels

A

No, they can only occupy the fixed orbits

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4
Q

How can electrons move between energy levels

A

They absorb or emit energy in the form of electromagnetic radiation or light

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5
Q

What does ‘n’ stand for

A

The electron shells of an atom

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6
Q

As the value of n increases the energy levels

A

Get closer together

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7
Q

What did the Niels Bohr model of the atom propose

A
  • Electrons revolve around the nucleus in circular orbits
  • The electron orbits correspond to certain energy levels or shells in the atom
  • Electrons can only occupy these fixed energy eves and cannot exist between two energy levels
  • Electron orbits of a larger radii correspond to a higher energy level
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8
Q

Which electrons are involved in chemical reactions

A

The valence electrons (outer most electrons)

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9
Q

What can you predict if you know the number of valence electrons in an atom

A

The chemical properties of an element

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10
Q

How are the elements in a periodic table arranged

A
  • In order of increasing atomic number (rows)
  • The number of electrons in the outer shell of an element (columns)
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11
Q

Groups

A
  • Going down in vertical columns
  • Number of valence electrons in the atom of an element
  • Labelled 1-18
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12
Q

Periods

A
  • Across the table in rows
  • Arrangement of electrons in each atom = to the number of occupied electron shells in the atom
  • Labelled 1-7
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13
Q

Helium

A
  • In group 18 but only has 2 valence electrons
  • it is placed here however as it is unreactive
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14
Q

What do the elements in the same group have in common

A

Chemical properties

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15
Q

Alkali metals properties

A
  • Highly reactive with water and oxygen
  • Relatively soft metals
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16
Q

Halogens properties

A
  • Coloured
  • Highly reactive
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17
Q

Noble gases properties

A
  • Stable
  • Inert/ low reactivity
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18
Q

Electron configuration

A

The number of electrons in each shell, separated by commas starting at the shell with the lowest energy

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19
Q

Ion

A

-A positively or negatively charged atom
- When protons and electrons don’t equal the same amount

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20
Q

How do ions form

A
  • Through chemical reactions or addition of energy the total number of electrons in an atom increases or decreases, this changes the electron configuration
  • The electrons are gained or lost from the valence shell as they are the most weakly held electrons
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21
Q

Negatively charge ions

A

Anions

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22
Q

Positively charged ions

A

Cations

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23
Q

Trends in the periodic table

A

Summarises the relative properties of elements and explains the trends observed in those properties

24
Q

List the trends in the periodic table

A
  • Electronegativity
  • Electrostatic attraction
  • Core charge
  • Atomic radius
  • Ionisation energy
  • Metallic character
25
Q

Electrostatic attraction

A

Holds individual atoms together, holds atoms together in molecules and molecules together in all forms of matter

26
Q

Coulomb’s Law

A

F = k x q1q2/r^2

  • q1 and q2 = the magnitudes of the charges involved
  • r = the distance between them
  • k = a constant
  • F = force of attraction
27
Q

Electrostatic attraction strength

A
  • If the size of the charge increases so does the electrostatic attraction
  • As the distance between the charges increase, the attraction decreases
  • The strength of the electrostatic attraction is directly proportional to the magnitude of the charges involved
28
Q

Core charge (effective nuclear charge)

A

A measure of the attractive force felt by the valence electrons towards the nucleus

29
Q

Core charge equation

A

Number of protons - the number of total inner shell electrons

30
Q

Core charge stays the same… and increases…

A
  • Down a group
  • Across a period (L-R)
31
Q

Nuclei- valence electron distance stays the same… and increases…

A
  • Across a period (L-R)
  • Down a group
32
Q

Atomic radius

A
  • Measurement for the size of atoms
  • From nucleus to valence electrons
33
Q

Atomic radius decreases

A

Across a period

34
Q

Atomic radius increases

A

Down a group

35
Q

Ionisation

A

The process of removing electrons from an atom to become an ion

36
Q

First ionisation energy

A

The energy which is required to remove the first valence electron from an atom in the gas phase

37
Q

Ionisation energy

A
  • The energy required to remove one electron from an atom of an element in the gas phase
  • The stronger the valence electron is held the more ionisation energy is needed
38
Q

First ionisation energy

A

The energy required to remove the first valence electron from an atom in the gas phase

39
Q

Ionisation energy increases

A

Across a period

40
Q

Ionisation energy decreases

A

Down a group

41
Q

Successive ionisation energy

A

The energy required to achieve the sequential removal of electrons from an atom

42
Q

Metallic character

A

How closely an element exhibits the properties commonly associated with metals (mostly that they lose an electron to form a cation)

43
Q

Electronegativity

A

The ability of an atom to attract electrons in a covalent bond towards itself

44
Q

Electronegativity increases

A

Across a period

45
Q

Electronegativity decreases

A

Down a group

46
Q

What can electronegativity predict

A

The type of bonding

47
Q

Ground state

A

Lowest energy level of the atom

48
Q

Excited state

A

The higher energy state of the atom

49
Q

How is absorbed energy emitted

A

Electromagnetic radiation

50
Q

Electromagnetic radiation examples

A
  • X-ray
  • Visible light
  • Radio waves
51
Q

What are different colours of lights

A

Electromagnetic radiation of varying wavelengths and frequencies

52
Q

Spectroscopy

A

The study of the interaction between matter and electromagnetic radiation

53
Q

Limitations of the flame test

A
  • Qualitative data only
  • Only a small range of metals are able to be detected
  • Metals in low concentrations can be difficult to observe
  • Mixtures of metals produce confusing results
54
Q

Atomic Absorption Spectroscopy

A

The direct relationship between the concentration of the sample being analysed and the light it absorbs

55
Q

Positives of the AAS

A
  • Provides both qualitative and quantitative data
  • More than 70 elements can be analysed
  • Can detect very low amounts of concentration
  • Highly selective > regularly used with mixtures

Chemistry (11 decks)

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