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Physchem > Atomic Theory Sample exam questions > Flashcards

Atomic Theory Sample exam questions Flashcards

1
Q

Define atomic orbital

A

Region of space around the nucleus where electrons are most likely to be found

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2
Q

Define first ionisation energy

A

The minimum energy required to remove most loosely bound electron from a neutral gaseous element in its ground state

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3
Q

Explain general increase in first ionisation energy across a period

A
  • Increases across
  • Decreasing atomic radius
  • Increase in effective nuclear charge
  • Screening effect remains the same
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4
Q

Define Atomic number

A

Number of protons in the nucleus of an atom

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5
Q

Define atomic mass

A

Number of protons and neutrons in the nucleus of an atom

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6
Q

Define Relative atomic mass

A

The average mass of an atom compared to 1/12th the mass of the carbon 12 isotope

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7
Q

Define electronegativity

A

The measure of the attraction that an atom has for the shared pair of an electron in a covalent bond

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8
Q

Describe how electronegativity values can be used to predict the type of bonds that occur between a pair of atoms

A

<0.4= pure covalent bond
>0.4 but <1.7=polar covalent bond
>1.7= ionic bond

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9
Q

explain the trend of atomic radius down the group

A
  • Increases
  • nuclear charge increases
  • screening effect increases
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10
Q

explain why alkali metals are so reactive

A
  • Group one
  • 1 electron in its outer energy level
  • it loses that one electron to follow the octet rule
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11
Q

Explain why alkali metals become more reactive down the group

A
  • atomic radius increases down the group
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12
Q

Explain why bromine is less reactive than chlorine

A
  • The electronegativity of Bromine is smaller

- The atomic radius is larger than chlorine

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13
Q

Explain why neon is chemically inert

A

8 electrons in its outside shell

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14
Q

How is an ionic bond formed between a pair of atoms

A
  • Electrons are transferred to form a positive ion

- To form a negative ion

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15
Q

How is a double covalent bond formed between a pair of atoms

A

The sharing of 2 pairs of electrons

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16
Q

transition metal

A

one that forms an ion with an incomplete d-subshell

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17
Q

Give 2 characteristic properties of transition metals

A

variable valency

acts as a catalyst

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18
Q

what is a crystal

A

solid with a repeating pattern of atoms

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19
Q

Describe the bonding in a metallic crystal

A

positive ions with valence electrons being shared

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20
Q

would you expect nickel to be a good electrical conductor, why?

A

yes

electrons are free to move

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21
Q

Account for the visible lines in the hydrogen emission spectrum

A
  • Electrons occupy the ground state
  • they move to excited energy levels due to heat or electricity
  • it’s unstable and falls back down
  • light is released
22
Q

Why is there a sharp increase from the 11th to the 12th ionisation energy

A

12th electron is taken from a new energy level which is closer to the positive nucleus

23
Q

Intermolecular forces

A

forces that exist between molecules

24
Q

Name the type of intermolecular force between

i) ammonia gas
ii) water
iii) methane gas

A

i) hydrogen bonds
ii) hydrogen bonds
iii) van der Waals forces

25
Q

explain in terms of intermolecular forces why the NH3 is very soluble in water

A
  • water and ammonia are both polar molecules
  • one side of each is slightly negative and the other side is positive
  • hydrogen bonds can form between them and ammonia dissolves
  • “like dissolves like”
26
Q

Relative atomic mass

A

The average mass of one element compared to 1/12th the carbon 12 isotope

27
Q

Why do electronegative values decrease across a period

A
  • nuclear charge increases
  • electrons enter the same energy level
  • screening effect stays the same
  • Electronegativity values decrease as the atomic radius increases
28
Q

Covalent crystal bonding

A

Each carbon atom forms four covalent bonds with four more carbon atoms

29
Q

What are the bonding forces in an ionic crystal

A

Ionic bonds

30
Q

What are bonding forces that hold molecular crystals together

A

Intermolecular forces

31
Q

Explain why the molecular crystals usually have lower melting points than ionic crystals

A

Intermolecular bonds in molecular crystals are weaker than ionic bonds which are easier to break

32
Q

What type of crystal is characterised by a solid having a lustre and very high mobility of electrons throughout the structure

A

Metallic

33
Q

Atomic Number

A

Number of protons in an atom of a nucleus

34
Q

First ionisation energy

A

The minimum energy required to remove the most loosely bound electron from a neutral gaseous element in its ground state

35
Q

A general increase in ionisation energy

A
  • Nuclear charge is increasing
  • Atomic radius is decreasing
  • Screening effect stays the same
36
Q

why is there a lot of energy required to ionise a neon atom

A

Neon is a noble gas
It has 8 electrons in its outer shell
It is stable and does not want to lose an electron

37
Q

Why is the first ionisation energy of xenon higher than neon

A

Xenon has a larger atomic radius than neon

38
Q

Define atomic orbital

A

A region in space around the nucleus where electrons are most likely to be found

39
Q

What’s the difference between the 2s and 2p orbitals of neon

A

S is lower in energy than the P orbital

Px, Py, Pz orbitals differ in direction

40
Q

Why do electron pairs repel each other

A

Like charges repel

Electrons are negatively charged so they repel each other at an energy level

41
Q

Compare the magnitude of repulsion between the two lone pairs and repulsion between two bonding pairs of electrons

A

Repulsion between two lone pairs of electrons is greater than the repulsion between two bond pairs
Lone pairs of electrons are closer to the nucleus than bond pairs

42
Q

Distinguish between ground and excited state

A

Ground: electrons occupy the lowest available energy level
Excited: Electrons occupy a higher energy level

43
Q

2 Ways an electron can become excited

A

Heat

Electricity

44
Q

Why is there no yellow line in the hydrogen emission spectrum

A

No transition equal in energy to produce a yellow line

Electrons cant occupy spaces between energy levels

45
Q

Describe how you carried out a flame test

A
  • Soak a wooden splint in water
  • dip in salt
  • place splint into the blue flame of a bunsen burner
46
Q

Define mass number

A

Number of neutrons and protons in an atom of a nucleus

47
Q

Define relative atomic mass

A

Mass of an atom of an element compared to 1/12th the carbon-12 isotope

48
Q

Define electronegativity

A

The measure of the force of attraction of an atom in a molecule for the shared pair of electrons in a covalent bond

49
Q

Identify a type of bond in a water molecule

A

Polar covalent bond

50
Q

Predict solubility of methane

A

Methane is not soluble in water
water is polar
methane is non-polar
Like dissolves like

Physchem (4 decks)

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