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Chem 150 > Exam 2 > Flashcards

Exam 2 Flashcards

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1
Q

Ionic bonding

A

electrostatic attraction between oppositely charged particles

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2
Q

Lattice energy

A

energy change required to form a mole of crystalline lattice from a mole of each gaseous ion

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3
Q

What happens as distance between ions increases?

A

Lattice energy decreases

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4
Q

Covalent bonding

A

involves the sharing of electrons between 2 non metallic elements

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5
Q

What type of bonding requires an empirical formula?

A

Covalent bonding

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6
Q

Law of multiple proportions

A

same elements, but different ratio equals a completely different ratio

refers to covalent bonds

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7
Q

Are covalent compounds molecules?

A

yes

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8
Q

Are ionic compounds molecules?

A

no

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9
Q

Lattice

A

positive and negative ions come together to form ionic compounds in a lattice structure

this structure consists of a positive cation surrounded by negative anions

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10
Q

Properties of the ionic lattice

A

since there are so many attractive forces, it takes a large amount of energy to break them apart

this gives ionic lattices a high melting point

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11
Q

Properties of covalent bonds

A

the covalent bonds inside the molecules are very strong

however, the forces attracting neighboring molecules are very weak

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12
Q

Which has a lower melting point molecular bonds or ionic bonds?

A

Molecular bonds

because of weak force of attraction between molecules

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13
Q

Is lattice energy positive or negative?

A

It depends on how you define it

If defined as splitting up a lattice into gaseous ions, then it will always be negative since it is an exothermic reaction (energy released)

If defined as a combining of gaseous ions into a lattice, then it will be positive since it is an endothermic reaction (energy required)

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14
Q

Relationship between lattice energy and ionic bond strength

A

Greater lattice energy = greater ionic bond strength

It takes more energy to break apart ions since there is less distance between them and greater force of attraction

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15
Q

What is lattice energy directly proportional to?

A

Charge

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16
Q

What is lattice energy indirectly proportional to?

A

Distance between ions

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17
Q

What happens to potential energy as two same charges get closer together?

A

Force of attraction decreases

So potential energy increases

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18
Q

What happens to potential energy as two opposite charges get closer together?

A

Force of attraction increases

So potential energy decreases

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19
Q

Potential energy “calibration” for opposite charges

A

Goes from more negative to less negative as distance between the electron and nucleus increases

as distance increases (force of attraction decreases), potential energy approaches 0 (increases)

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20
Q

Is force of attraction greater in He+ or H (have the same protons)?

A

Greater in He+ because helium has more protons than hydrogen pulling at the one electron

21
Q

Why would Cr have the configuration [Ar] 4s1 3d5?

A

Having a two paired electrons in the 4s orbital makes electron-electron repulsion increase and thus potential energy goes up

Therefore, want to have no paired electrons because this decreases electron-electron repulsion and thus gives the overall system less energy

22
Q

Effective nuclear charge

A

the actual magnitude of the positive charge experienced by an electron

= (# of protons) - core electron shielding

23
Q

What is the relationship between effective nuclear charge and electron shielding?

A

As electron shielding increases, effective nuclear charge decreases

24
Q

Why are orbitals with the same energy level degenerate in hydrogen but not in multi-electron atoms?

A

Since hydrogen only has 1 electron, there are not other atoms to reduce the effective nuclear charge

For example, 2p is being affected by shielding from 2s so it bumps up to a higher energy level (feels less force of attraction, so has higher energy)

25
Q

Relationship between effective nuclear charge and atomic radius

A

As effective nuclear charge increases, atomic radius decreases since there is a stronger force of attraction

26
Q

What happens to atomic radius as you go down? Explain.

A

Increases

Since the effective nuclear charge increases, but the core electrons also increase, there is less force of attraction, so the radius increases

Also, n is increasing which indicates bigger size because there are more shells which increases shielding and distance

27
Q

Why does atomic radius decrease as you go across?

A

Since protons increase but core electrons (in the same shell) remain the same, there is a net effective nuclear charge increase which pulls electrons closer, thus decreasing the radius

28
Q

Ionization energy

A

the amount of energy needed to remove one mole of electrons from one mole of a certain element

29
Q

Why does it take energy to remove an electron from an atom?

A

Need to overcome the force of attraction

Also think, you are moving to n = ∞, so need to absorb energy to increase n-level

30
Q

What happens as to ionization energy as you go down a column? Explain.

A

Ionization energy decreases

Since n-level increases, the atomic radius increases and there is less force of attraction. Therefore, it takes less energy to remove an electron

31
Q

What happens to ionization energy as you go across a column? Explain.

A

Ionization energy increases

Since effective nuclear charge increases (more protons), force of attraction is stronger and it takes more energy to remove an electron

32
Q

Why is it easier to remove electrons from the p-orbital than the s-orbital? (lower ionization energy)

A

s-orbital is more likely to be closer to the nucleus which results in stronger force of attraction than p-orbital

33
Q

Why does oxygen have lower ionization energy than nitrogen?

A

Oxygen has a paired electron orbital, whereas nitrogen has all unpaired electrons

Since paired electron orbitals create electron-electron repulsion and lower effective nuclear charge, it is easier to remove an electron since less force of attraction

34
Q

Electron shielding

A

describes the decrease in force of attraction between the outermost electron and the nucleus due to paired electrons (greater electron-electron repulsion)

35
Q

Describe electron shielding in terms of energy

A

Since paired electron orbitals create electron-electron repulsion, there is an increase in energy

an increase in energy indicates a decrease in force of attraction

therefore, there is less of a force of attraction due to electron shielding

36
Q

Atomic number (Z)

A

number of protons

the total positive charge in an atom

37
Q

Why is a cation smaller than its corresponding atom?

A

Since a cation looses an electron, each remaining electron experiences a greater effective nuclear charge from the protons

therefore, force of attraction increases and cations are smaller

38
Q

Why is an anion larger than its corresponding atom?

A

Since an anion gains an electron, each electron experiences a lesser effective nuclear charge from the protons

therefore, force of attraction decreases and anions are larger

39
Q

Why are electrons for transition metal cations removed from the s subshell instead of the d subshell?

A

3d is more likely to be closer to the nucleus based on radial probability, thus giving it a stronger force of attraction than 4s. Therefore, electrons are removed from 4s

40
Q

f-block and shielding

A

the f-block are poor shielders

this creates a big increase in protons as we move down the periodic table that is unaccompanied by electrons not shielding which equals an overall higher effective nuclear charge

41
Q

Why is columbs law (force) directly proportional to lattice energy?

A

Because to break apart a strong force of attraction it requires a greater input of energy

42
Q

Why does the atomic radius not increase as you go across the transition metals?

A

As you add electrons to the d-block, you also increase shielding and therefore there is no net change in effective nuclear charge

43
Q

Nobel peace prize winners in chem 2018

A

George Smith, Frances Arnold, Greg Winter

44
Q

What to use when comparing force of attraction using Coulomb’s law?

A

Magnitude

45
Q

What to use when comparing lattice energies?

A

Magnitude

46
Q

Relationship between ionic radius and lattice energy

A

Ionic radius is inversely proportional to lattice energy

47
Q

Why do ionic compounds want to form?

A

Energetically favorable

In addition, being isoelectronic with noble gas configuration provides higher Zeff and correspondingly lower potential energy

48
Q

Why is energy released when a lattice forms?

A

Phase change from gas to solid (higher energy to lower energy)

Also, force of attraction increases, so potential energy decreases

49
Q

What happens as the electron moves closer to another ion in order to form a lattice?

A

Force of attraction increases

Potential energy decreases

Chem 150 (10 decks)

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