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Uncategorized > General Chemistry I > Flashcards

General Chemistry I Flashcards

1
Q

In the Bohr Model, what does the Hydrogen electron orbit?

A

Nucleus

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2
Q

In quantum mechanics, where does the Hydrogen electron exist?

A

In a spherical probability cloud around the nucleus.

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3
Q

In quantum mechanics, what does the principle quantum number, n, define?

A

The principle quantum number, n, defines what shell the electron is in.

n values start from one: 1,2,3 …etc.

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4
Q

As the principle quantum number, n, increases, the energy [increases/decreases].

A

increases

Only if subshells are the same.

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5
Q

How many electrons per orbital?

A

Two electrons.

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6
Q

How many orbitals are there per shell?

A

n2

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7
Q

How many electrons per shell?

A

2n2

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8
Q

Electrons are naturally found in which state?

A

Ground State

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9
Q

To which state does an electron move when it absorbs energy?

A

Excited State

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10
Q

Which is higher in energy, the excited states or the ground state?

A

Excited States

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11
Q

Energy is ______ when a electron moves from excited to ground state.

A

released

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12
Q

Define:

Absorption Spectrum

A

The electromagnetic spectrum, broken by a specific pattern of dark lines or bands, observed when radiation traverses a particular absorbing medium. The absorption pattern is unique and can be used to identify the material.

Example of Chlorophyll Absorption Spectrum

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13
Q

Define:

Emission Spectrum

A

The spectrum of bright lines, bands, or continuous radiation characteristic of and determined by a specific emitting substance subjected to a specific kind of excitation.

Emission Spectrum of Different Elements

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14
Q

In quantum mechanics, how many orbitals can be found in the s shell?

A

One orbital.

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15
Q

In quantum mechanics, how many orbitals can be found in the p shell?

A

Three orbitals.

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16
Q

In quantum mechanics, how many orbitals can be found in the d shell?

A

Five orbitals.

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17
Q

In quantum mechanics, how many orbitals can be found in the f shell?

A

Seven orbitals.

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18
Q

In quantum mechanics, what does l symbolize?

A

l is the angular momentum quantum number. l ranges from 0 to n-1.

l=0,1,2,3 for s,p,d,f respectively.

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19
Q

In quantum mechanics,what do the letters s,p,d,f symbolize?

A

The letters s,p,d,f symbolize the subshells.

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20
Q

In quantum mechanics, what is the maximum number of electrons found in an orbital?

A

Two electrons

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21
Q

In quantum mechanics, for a given shell, higher subshells have [higher/lower] energy.

A

higher

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22
Q

In quantum mechanics,what does the letter m signify?

A

The letter m signify’s the magnetic quantum number.

m ranges from -1 to 1, including 0.

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23
Q

In quantum mechanics, what does the letter s signify?

A

The letter s signify’s the spin quantum number.

s is either -1/2 or 1/2.

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24
Q

What are the common names and geometric shapes for orbitals s, p, d?

A
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25
Q

In quantum mechanics, subshells are filled by placing electrons in the _____ subshell first.

A

lowest

26
Q

Arrange the following subshells in increasing energy: 4s, 3d, 5s, 2s, 4d, 5p, 1s, 6s, 4f, 2p, 5d, 6p, 3s, 7s, 5f, 6d, 3p, 4p

A

Subshells arranged in increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d

27
Q

Write the orbital diagram for the element Flourine.

A
28
Q

In quantum mechanics, what is the conventional notation for electronic structure?

A
29
Q

Define:

Aufbau’s Principle

A

Aufbau’s Principle:

shells / subshells of lower energy get filled first.

For example, 1s fills first, then 2s, then 2p …etc. Review the exact order of energies because later on, the d subshells get filled after the s.

30
Q

Define:

Hund’s Rule

A

Hund’s Rule:

when you fill a subshell with more than 1 orbital (p, d, f), you first fill each orbital with a single electron and with the same spin

The reason for Hund’s rule is that electron-electron repulsion in doubly occupied orbitals make them higher in energy than singly occupied orbitals.

31
Q

Define:

Pauli’s Exclusion Principle

A

Pauli Exclusion Principle:

Two electrons in the same orbital must be of different spins.

32
Q

What is the equation to for effective nuclear charge in terms of shielding electrons?

A

Effective nuclear charge = nuclear charge - shielding electrons.

33
Q

The higher the effective nuclear charge for an electron, the _____ the ionization energy.

A

higher

Effective nuclear charge increases for outer electrons as you go across (left to right) the periodic table.

34
Q

Alkaline earth metals have [number] valence electrons and [high/low] ionizations energy.

A

Alkaline earth metals have_ two _valence electrons and low ionizations energy.

Alkaline earth metals are quite reactive and strive to lose both electrons to achieve an empty valence shell.
Alkaline earth metals are more reactive as you go down because of increasing radii.

Most commonly found in the +2 oxidation state.

35
Q

Halogens have [number] valence electrons and [high/low] electron affinity.

A

Halogens have seven valence electrons and high electron affinity.

Halogens are very reactive and strive to gain one electron to achieve full valence shell.
Halogens are more reactive as you go up a column because of decreasing radii.

Most commonly found in the -1 oxidation state.

36
Q

Noble Gases have [number] valence electrons and [high/low] electron affinity and [high/low] ionization energy.

A

Noble Gases have eight valence electrons and low electron affinity and high ionization energy.

Noble gases are rarely react.

Found in the oxidation state of 0.

37
Q

Alkali Metals have [number] valence electrons and [high/low] ionization energy.

A

Alkali Metals have one valence electrons and low ionization energy.

Alkaline metals are very reactive and strive to lose that electron to achieve an empty valence shell. Alkaline metals are more reactive as you go down because of increasing radii.

Most commonly found in the +1 oxidation state.

38
Q

Why do transition metals have high conductivity?

A

Transition metals have a “sea of mobile electrons”: free flowing (loosely bound) outer d electrons.

In the presence of ligands (when in a chemical complex), the d orbitals become nondegenerate (different in energy).
Electron transitions between nondegenerate d orbitals gives transition metal complexes vivid colors.Varied oxidation states - but always +.

39
Q

Define:

Representative Element

A

Representative elements include the s block and the p block of the periodic table because they have no free flowing (loosely bound) outer d electrons.

Valence shell fills from left (1 electron) to right (8 electrons).Standard nomenclature from left to right: I A, II A, III A, IV A, V A, VI A, VII A, VIII A.

40
Q

What are the main chemical differences between metals and nonmetals?

A

Metals

  1. Likes to lose electrons to gain a + oxidation state (good reducing agent).
  2. Lower electronegativity - partially positive in a covalent bond with non-metal.
  3. Forms basic oxides.

Non-metals

  1. Likes to gain electrons to form a - oxidation state (good oxidizing agent).
  2. Higher electronegativity - partially negative in a covalent bond with metal.
  3. Forms acidic oxides.
41
Q

Define:

Oxygen Group

A

The group (column) that contains oxygen.

42
Q

What are the main physical differences between metals and nonmetals?

A

Metals

  1. Good conductor of heat and electricity
  2. Malleable, ductile, luster, solid at room temp(except Hg)

Non-metals

  1. Poor conductor of heat and electricity
  2. Solid, liquid, or gas at room temp. Brittle if solid and without luster.
43
Q

Which element is the most similar to Oxygen?

A

Sulfur

44
Q

Describe the general trend of electron affinity on the periodic table.

A

As you go down a group, electron affinity decreases because of larger radii.As you go across (left to right) a row, electron affinity increases.
Highest peaks are for the halogens.Lowest for noble gases.
Local minima occurs for filled subshells and half-filled p subshells.

45
Q

Which quantum number defines electron shell?

A

n

46
Q

As you move from Ne to Na, what happens to the effective nuclear charge?

A

The effective nuclear charge decreases from Ne to Na due to shielding.

47
Q

Size [increase/decrease] as you go down a column.

Size [increase/decrease] as you go across (to the right of) a row.

A

Size increases as you go down a column.

Size decreases as you go across (to the right of) a row

48
Q

Define:

Electronegativity

A

A chemical property that describes the ability of an atom (or, more rarely, a functional group) to attract electrons (or electron density) towards itself.

49
Q

Which element is the most electronegative?

A

Fluorine

50
Q

Define:

Electron Affinity

A

Electron Affinity is the energy change when an electron is added to the neutral species to form a negative ion.

51
Q

In quantum mechanics, how many electrons can be found in the s shell? p? d? f?

A

s subshells hold 1x2=2 electrons,

p subshells holds 3x2=6 electrons,

d subshells holds 5x2=10 electrons,

f subshells holds 7x2=14 electrons.

52
Q

Describe the pattern of the periodic table in terms of subshells and energy.

A

Starting from the first row, going across, both hydrogen and helium is 1s. Next row: 2s then 2p. Third row: 3s then 3p. Fourth row: 4s, 3d, then 4p. Fifth row: 5s, 4d, then 5p. Sixth row: 6s, 4f, 5d, then 6p. Last row: 7s, 5f, then 6d. The pattern we get from looking at the periodic table is exactly in the order of increasing energy.

53
Q

In quantum mechanics, what do the columns represent on the periodic table?

A

For a given subshell, the columns represent how many electrons are in that subshell. For example, the fifth column of the d subshells contain elements that have 5 electrons in that subshell.

The number of columns for each subshell indicate the maximum number of electrons that subshell can hold. For example, the d subshells have 10 columns showing that d orbitals can hold 10 electrons total.

54
Q

In quantum mechanics, what is different about d4 and d9 elements when writing the subshell notation?

A

Watch out for d4 and d9 elements. Instead of s2d4, it’s s1d5 and s1d10 because they want to achieve a half-full or full d subshell.

55
Q

Define:

Shielding Electrons

A

The electrons in the energy levels between the nucleus and the valence electrons.

MCAT questions usually give you a diagram of the Bohr model, in which case, shielding electrons are those that orbits at a smaller radius.

56
Q

Describe the classification of elements into groups by electronic structure, for the periodic table.

A

The periodic table: classification of elements into groups by electronic structure; physical and chemical properties of elements.

57
Q

Define:

First Ionization Energy

A

The energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

58
Q

Define:

Second Ionization Energy

A

The energy required to remove a second valence electron from the univalent ion to form the divalent ion.

59
Q

Describe the general trend of ionization energy on the periodic table.

A

Ionization energy decreases as you go down because of increasing radii.Ionization energy increases as you go right because of decreasing radii.

Highest peaks are noble gases.Lowest troughs are alkali metals.
Local maxima occurs for filled subshells and half-filled p subshells.Second ionization energy is always higher than the first ionization energy (usually a lot higher).
Alkali metals and hydrogen: first ionization energy very low. Second ionization much higher.Alkaline earth metals: first ionization energy low. Second ionization energy also low.

60
Q

Define:

Electron Affinity

A

Electron affinity is the amount of energy released when an atom gains an electron.

61
Q

Describe the general trend of electronegativity on the periodic table.

A

Electronegativity increases toward the top right.
Things around fluorine are highly electronegative: N, O, F, Cl, Br.Halogens are electronegative, especially toward the top of the group.
Noble gases can be very electronegative if they participate in bond formation (Kr and Xe).Non-metals are more electronegative than metals.
Covalent bond is a sharing of electrons between elements.The more electronegative element in a covalent bond gets a larger share of the electrons and has a partial negative charge
The less electronegative (more electropositive) element in a covalent bond gets a smaller share of the electrons and has a partial positive charge.If the electronegativity difference is too great, an ionic bond occurs instead of a covalent one.
Ionic bonds result from a complete transfer of electrons from the electropositive element to the electronegative element.

62
Q

Describe the general trend of atomic radii on the periodic table.

A

Size increases as you go down a column and decreases as you go across (to the right of) a row.
Atomic sizes may overlap if you zigzag on the periodic table.

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