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BTEC Unit 1 > Chemistry > Flashcards

Chemistry Flashcards

1
Q

Name the three sub orbitals (sub shells)

A

s, p, d

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2
Q

What sub orbitals can be found on the first energy level and how many?

A

1 x s orbital

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3
Q

What sub orbitals can be found on the second energy level and how many?

A

1 x s and 3 x p

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4
Q

What sub orbitals can be found on the third energy level and how many?

A

1 x s, 3 x p and 5 x d

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5
Q

How many electrons can be placed in each orbital?

A

Two

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6
Q

How do electrons fill the 3 p orbitals?

A

Singularly first (to avoid repelling) and then they pair

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7
Q

Write the electron configuration for phosphorus

A

1s2 2s2 2p6 3s2 3p3

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8
Q

Write the electron configuration for chromium

A

1s2 2s2 2p6 3s2 3p6 4s1 3d5 (half filled orbitals)

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9
Q

Why do electrons fill the 4s before the 3d?

A

It is lower in energy

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10
Q

Explain the arrangement of 2p electrons for oxygen

A

2 electrons in Px and 1 each in Py and Pz. This happens because electrons fill orbitals singularly and then pair

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11
Q

Is potassium in the s, p or d block?

A

s block

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12
Q

Is fluorine in the s, p or d block?

A

p block

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13
Q

Is titanium in the s, p or d block?

A

d block

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14
Q

What is an ion?

A

An atom that has lost or gained electrons

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15
Q

What ions do group 2 elements make?

A

+2 (lose 2 electrons)

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16
Q

What ions do group 5 elements make?

A

-3 (gains 3 electrons)

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17
Q

Write the electron configuration for Ti 2+

A

1s2 2s2 2p6 3s2 3p6 3d2 (empty 4s first)

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18
Q

What is the formula for a nitrate ion?

A

NO3-

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19
Q

What is the formula for a sulphate ion?

A

SO42-

20
Q

What is the formula for a hydroxide ion?

A

OH-

21
Q

Between which atoms does ionic bonding occur?

A

Metal and non-metal

22
Q

What happens to the electrons in ionic bonding?

A

They are transferred between atoms forming ions

23
Q

What atoms form covalent bonds?

A

Non metals

24
Q

What happens to electrons in a covalent bond?

A

They are shared

25
Q

What happens to electrons in metallic bonding?

A

They are delocalised and can move freely within the structure

26
Q

What is the nature of the bonding in ionic compounds?

A

electrostatic attraction between oppositely charged ions

27
Q

What is the nature of the bonding in colavent compounds?

A

Electrostatic attraction between positively charged nucleus and negatively charged electrons

28
Q

Describe the structure of an ionic compound

A

A giant lattice of oppositiely charged ions

29
Q

Describe the structure of a metal

A

A giant lattice of positively charged metals ions with delocalised electrons

30
Q

Why do ionic compounds have high melting points?

A

Strong electrostatic forces of attraction that require lots of energy to break

31
Q

Why do metals have high melting points?

A

Strong electrostatic force that requires lots of energy to break

32
Q

Why do simple covalent compounds such as O2, CO2 and H2O have low boiling points?

A

Weak intermolecular forces between molecules that require little energy to break

33
Q

Why do metals conduct electricity?

A

They have delocalise electrons that are free to move and carry charge

34
Q

WHy do ionic compounds conduct electricity when molten/dissolved?

A

Ions are free to move and carry charge

35
Q

Why do ionic compounds not conduct electricity when solid?

A

Their ions are not free to move and carry charge

36
Q

Define electronegativity

A

The power of an atom to attract electrons within a covalent bond

37
Q

Which are the most electronegative atoms?

A

F, O, N and Cl (top right corner)

38
Q

How does the electronegativity change down the group?

A

Decreases - outer electron further from the nucleus (more shells) and more shielded. Therefore weaker attraction

39
Q

How does the electronegativity change across a period?

A

Increases because there is a higher effective charge in the nucleus therefore stronger attraction for electron

40
Q

How does atomic (ionic) radii change down a group?

A

Increases. Inctrease number of shells and increased shielding. Therefore weaker attraction for outer electron

41
Q

Which three factors affect atomic radii?

A

Nuclear charge, number shells and shielding

42
Q

How does atomic (ionic) radii change across a period?

A

Decreases - effective nuclear charge increases and therefore attraction for outer electron is stronger

43
Q

Define first ionisation energy

A

The energy required to remove one mole of electrons from one mole of gaseous atoms

M –> M+ + e

44
Q

Define first electron affinity

A

The energy released when one mole of gaseous atoms gains one mole of electrons

M + e –> M-

45
Q

Describe the general trend of ionisation energy across a period

Why are there exceptions?

A

Increases as the size of the nuclear charge increases and the electrostatic attraction is stronger meaning more energy is required to remove an electron.

Two exceptions (the first is because 3p is higher in energy than 3s and the second where the electron is being removed from 3p4)

BTEC Unit 1 (2 decks)

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