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Alevel Chemistry > Atomic Structure,Bonding, energetics > Flashcards

Atomic Structure,Bonding, energetics Flashcards

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1
Q

Isotopes

A

Isotopes are atoms of the same element with a different no. of neutrons

  • chemical properties of isotopes are similar
  • physical properties are different (due to mass - changes BP & MP)
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2
Q

proof of electron orbitals

A
  • Drop of ionisation energy between G5 & G6 is proof of electron orbitals.
  • Oxygen has a lower IE than nitrogen
  • There’s no difference between atomic radius & shielding, but there is more repulsion between a pair of electrons within an orbital so electrons in shared orbitals are easier to move.
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3
Q

Time Of Flight (TOF) mass spectrometry processs

A

1 - IONISATION
↳ vaporised gas is ionised via electrospray or electron
impact
2 - ACCELERATION
↳ The ions are accelerated by an electric field. (relies on particles being charged)
↳ all ions will have the same KE (velocity dependent on
mass)
3 - ION DRIFT
↳ ions drift through the chamber at different speeds under no electric field
4 - DETECTION
↳ Ions are detected by producing current on detection plate which creates a mass spectrum.

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4
Q

Trend in Ionisation energy across a period

A
  • ↑Nuclear charge across period ∵ increase in atomic no.
  • ↓ Atomic Radius
    ↳ atomic radius decreases slightly due to greater nuclear charge that pulls outer electrons closer
  • NO change in shielding effect ∵ same no. of electron shells
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5
Q

Trend in Ionisation energy across a group

A
  • ↑ Atomic Radius
    ↳ greater distance between nucleus and outermost electron, decreased attraction.
  • ↑ Shielding effect
    ↳ greater no. of inner shells, ∴ greater repulsion so decreased attraction.
  • negligible nuclear charge increase
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6
Q

Factors affecting IE

A

Shielding -
The repulsion of an electron from inner shell electrons.
- more inner shells present, lower IE
Atomic Radius -
↑ atomic radius, lower IE, outer electron further from nucleus so weaker attractive pull from nucleus.
Nuclear Charge -
higher nuclear charge the higher the ionisation energy, stronger attraction to outer electrons.

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7
Q

How is Electrospray Ionisation done?

A

Sample is dissolved in solvent
- High voltage is applied
- High voltage removes proton from solvent and attaches it to sample
- X(g) —> XH⁺(g)

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8
Q

How is Electron Impact Ionisation performed

A
  • Sample is vaporised
  • Hit with electrons from electron gun
  • The electrons knock off electrons from the molecule
  • X —-> X⁺+ e⁻

(often causes fragmentation)

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9
Q

Electron Configuration of Transition Metals

A

1s²2s²2p⁶3s²3p⁶4s¹3d⁵ <— Chromium
1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰ <—- Copper

+ When forming ions they lose their 4s electron first

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10
Q

Proof of electron sub-shells

A
  • Drop of ionisation energy between G2 & G3 is proof of sub-shells.
  • Beryllium has a greater ionisation energy than boron.
  • Boron’s outer electron occupies a 2p sub-shell which increases atomic radius & it’s electron experiences more shielding from inner electrons ( both 1s² and 2s²electrons shield it)
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11
Q

Properties of Water

A
  • Significantly higher boiling point than other hydrides
  • Surface tension ( how strongly molecules are held to the structure of a liquid)
    • H bonds exert a downwards force
  • High Viscosity
  • Density of ice lower than water because it freezes into regular lattice
    structure that are further apart so more open structure =
    less dense
  • solvent can dissolve simple alcohols and ionic compounds
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12
Q

Physical Properties of Metallic Structures

A
  • Metals have high boiling & Melting points
  • More delocalised electrons present, the higher the melting point of the metal.
    • because greater electrostatic attraction so stronger bond
  • Good conductors of heat & electricity because of delocalised electronns
  • insoluble unless liquid metal.
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13
Q

Charge Clouds

A
  • region where there is a high chance of an electron pair being present. / region of negative charge.
  • Can contain either boning pairs or lone pairs
  • Lone pair electron charge clouds repel more than bonding pair clouds
    Great repulsion LONE - LONE
    LONE - BOND
    BOND - BOND
    Lowest repulsion

<- VSEPR theory

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14
Q

Formation of temporary dipoles

A

Temporary dipoles are week intermolecular forces
- When two atoms come towards each other the electron CLOUDS repel each other
- causes sudden displacement of electrons to one side resulting in a temporary dipole.
- dipoles are constantly being created or destroyed.

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15
Q

ALL shapes of molecules: name,shapes & angles

A
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16
Q

Factors affecting the strength of a metallic bond

A
  • size of positive charge on the ions
    • greater charge =greater attraction to delocalised electrons
  • Size of Metal Ion
    • strength of metallic bond increases with decreasing metal
      ion size
    • shorter distance between positive nucleus and delocalised
      electrons.
  • Number of mobile electrons per atom
    • more mobile electrons contributing to the sea of electrons
      holding lattice together
17
Q

Graphite

A
  • Organised sheets of hexagons held together by Van Der Waals inbetween them
  • the layers can easily slide over each other because there is only weak intermolecular forces holding them together.

-good dry lubricant

18
Q

Bond Enthalpy

A
  • Energy required to break one mole of a particular covalent
    bond in gaseous state
  • The Bond Energy depends on Bond Length
    • the shorter the bond length, the stronger the bond . thus
      more energy is needed to break it.
  • Lower Bond Energy –> more reactive
19
Q

Hydrogen bonding

A

Polar bond - H” F,N,O’
- Hydrogen has a high charge density (only 1 shell) so the “ hydrogen can be attracted by lone pairs of nearby N,O,F molecules

+ due to hydrogen bonding it means HF is the hydrogen halide with the highest boiling point.

20
Q

Explain why the third ionisation energy of magnesium is much higher than the second ionisation energy of magnesium.

A

The electron is removed from lower energy level
- electron being removed is less shielded so more energy is needed to overcome attraction.

21
Q

Enthalpy of Formation (ΔH°f)

A

The enthalpy change when one mole of a substance is produced from its elements under standard conditions.

22
Q

Hess’s Law

A

The Enthalpy change is independent of the route taken.

23
Q

Enthalpy of Combustion (ΔH°c)

A

The enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions.

24
Q

Bond Enthalpy

A

The energy required to break one mole of the stated bond in a gaseous state, under standard conditions.

25
Q

How is the relative abundance of a molecule determined using a TOF mass spectrometer

A

At he detector/negative plate the ions GAIN an electron, causing a current
The relative abundance depends on the SIZE of the current.

26
Q

Why do isotopes have the same chemical properties?

A

Same electron configuration so there is no change in chemical properties

27
Q

Relative Atomic Mass

A

Average/mean of 1 atom of an element 1/12 mass of one atom of carbon-12

28
Q

Which is larger Na+ or F- and why?

A
  • F-
  • both Na+ and F- have same electron arrangement
  • sodium ion has more protons so attracts outer electrons closer/stronger attraction to outer electrons.
29
Q

Why does SiO2 have a larger boiling point than P4O10?

A

-SiO2 is macromolecular/ giant covalent
-therefore it has strong covalent bonds which require a lot of energy to overcome.
-P4O10 is simple molecular so only has Vander Waals forces acting between the molecules.

30
Q

Define enthalpy change

A

The change in heat energy during a reaction at constant pressure.

31
Q

why may a bond not have a mean bond enthalpy

A

it may be the only bond in the substance/reaction

32
Q

why may the experimental value for enthalpy of combustion be less than the data book value?

A
  • heat loss to the surroundings
    -incomplete combustion
33
Q

Standard enthalpy of combustion

A

The enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions.

34
Q

Why many a calorimeter heat change calculation not equal the enthalpy change?

A

the pressure in a calorimeter is not constant

35
Q

find the heat energy change using calorimeters in an efficient way

A
  • insulate beaker using a polystyrene cup to reduce heat loss
  • record the temperature by for a suitable time before adding the substance (if a solution) to establish an accurate initial temperature
  • record the temperature values at regular time intervals and plot the temperature results against time on a graph
  • extrapolate the graph to find your theoretical maximum temperature

Alevel Chemistry (11 decks)

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